MEDICAL 


ollege  of  Pharmacy 


California  College  of  Pharmacy 


QUALITATIVE 

CHEMICAL    ANALYSIS 


A  GUIDE  IN  QUALITATIVE  WORK,  WITH  DATA 

FOR    ANALYTICAL    OPERATIONS    AND1 

LABORATORY  METHODS 


BY 

ALBERT  B.  PRESCOTT 

AND 

OTIS  C.  JOHNSON 

PROFESSORS  IN    THE  UNIVERSITY   OF  MICHIGAN 


SEVENTH  EDITION,   THOROUGHLY  REVISED 


BY 

JOHN  C.  OLSEN,  A.M.,  PH.D. 

Professor  of  Chemicai  Engineering,  Polytechnic  Institute,  Brooklyn,  N.  Y. 

Author     of    "Quantitative     Chemical    Analysis" 

Editor,  Van  Nostrand's  "Chemical  Annual" 

California  College  of  Pharmacy 


NEW  YORK 

D.  VAN  NOSTRAND  COMPANY 

Eight  Warren  Street 

1920 


Copyright,  1916,  1917 
BY  D.  VAN  NOSTRAND  COMPANY 


Lo 


PREFACE  TO  THE  SEVENTH  EDITION. 


Since  the  last  revision  of  this  very  comprehensive  text  on  Qualitative 
Analysis  was  issued,  a  considerable  number  of  new  methods  of  analysis 
have  been  published,  and  some  of  these  have  been  found  to  be  valuable 
and  have  come  into  more  or  less  general  use. 

Many  of  the  new  methods  have  been  tried  out  and  some  have  been 
found  to  be  unreliable  unless  extreme  care  is  taken  or  only  pure  solu- 
tions used.  The  attempt  has  been  made  to  include  in  this  revision  only 
those  methods  which  have  been  found  to  be  reliable. 

A  very  marked  change  has  also  taken  place  in  the  method  of  presen- 
tation of  chemical  reactions.  Ionic  reactions  are  given  in  many  texts  to 
the  exclusion  of  molecular  reactions.  Physical  chemical  theories  are  also 
frequently  presented  and  discussed  at  great  length. 

The  attempt  has  been  made  in  the  revision  of  this  text  to  present 
only  briefly  the  modern  conceptions  of  solution,  leaving  a  fuller  presenta- 
tion for  the  lecturer  or  separate  texts  on  Physical  Chemistry.  Mole- 
cular reactions  have  also  been  largely  retained  in  the  belief  that  the 
material  to  be  analyzed  by  the  chemist  in  practical  work  is  quite  as 
often  a  molecular  compound  as  an  ion. 

All  molecular  weights,  solubilities  and  other  constants  of  the  elements 
and  their  principal  compounds  have  been  brought  up  to  date.  The 
principal  minerals,  methods  of  preparation  and  determinations  of  ele- 
ments, as  well  as  the  reactions,  have  been  revised,  references  being  given 
to  the  literature  as  heretofore. 

In  general,  the  attempt  has  been  made  to  retain  the  excellent  features 
of  this  text  which  have  given  it  such  an  extended  use  in  the  past,  both  as 
a  class  room  and  as  a  reference  text,  while  adding  the  valuable  results  of 
recent  progress  in  the  science. 

Acknowledgment  is  made  with  pleasure  of  valuable  assistance  ren- 
dered the  undersigned  in  preparing  the  revision  to  Professor  R.  J. 


PREFACE. 

Colony  of  Cooper  Union,  who  revised  the  sections  on  the  properties,  oc- 
currence and  preparation  of  the  metals ;  to  Mr.  Wm.  H.  Ulrich,  who  tried 
out  many  of  the  new  methods;  to  Mr.  N.  F.  Borg,  who  contributed  much 
of  the  revision  of  the  acids ;  and  to  Mr.  M.  P.  Matthias,  who  has  revised 
the  Index. 

J.  C.  OLSEN 
October  2,  1916. 


PREFACE  TO  THE  FIFTH  EDITION. 


In  this,  the  fifth  full  revision  of  this  manual,  the  text  has  been  re- 
written and  the  order  of  statement  in  good  part  recast.  The  subject- 
matter  is  enlarged  by  fully  one-half,  though  but  one  hundred  pages 
have  been  added  to  the  book. 

It  has  been  our  aim  to  bring  the  varied  resources  of  analysis  within 
reach,  placing  in  order  before  the  worker  the  leading  characteristics  of 
elements,  upon  the  relations  of  which  every  scheme  of  separation  de- 
pends. This  is  desired  for  the  working  chemist,  and  no  less  for  the 
working  student.  However  limited  may  be  the  range  of  his  work,  we 
would  not  contract  his  view  to  a  single  routine.  It  is  while  in  the 
course  of  qualitative  analysis  especially  that  the  student  is  forming 
his  personal  acquaintance  with  the  facts  of  chemical  change,  and  it  is 
not  well  that  his  outlook  should  be  cut  off  by  narrow  routine  at  this 
time. 

The  introductory  pages  upon  Operations  of  Analysis,  setting  forth 
some  of  the  foundations  of  qualitative  chemistry,  consist  of  matter 
restored  and  revised  from  the  editions  of  1874  and  1880.  This  sub- 
ject-matter, omitted  in  1888,  is  now  desired  by  teachers.  For  the  portion 
upon  Solution  and  lonization,  we  are  indebted  to  Dr.  Eugene  C.  Sulli- 
van, a  pupil  of  Professor  Ostwald,  now  teaching  qualitative  analysis. 
The  pages  upon  the  Periodic  System  have  been  added  to  afford  a  more 
connected  comparison  of  the  elements  than  that  undertaken  in  each 
group  by  itself,  in  previous  editions,  and  referred  to  in  the  preface  in 
1874.  The  use  of  notation  with  negative  bonds,  in  balancing  equations 
for  changes  of  oxidation,  introduced  by  one  of  the  authors  in  1880, 
has  been  retained  substantially  as  in  the  last  edition.  Other  authors 
adopt  the  same  notation  with  various  modifications.  For  the  present' 
revision  there  has  been  a  general  search  of  literature,  and  authorities 
are  given  for  what  is  less  commonly  known  or  more  deserving  of  further 

v 


VI  PREFACE. 

inquiry.  The  number  of  citations  is  so  large  that  to  save  room  special 
abbreviation  is  resorted  to. 

For  convenient  reference,  on  the  part  of  teachers,  students  and 
analysts  using  the  book,  the  section  for  each  element  and  each  acid  is 
arranged  in  uniform  divisions.  For  instance,  in  each  section,  solu- 
bilities are  given  in  paragraph  5,  the  action  of  alkalis  in  paragraph  6a, 
the  action  of  sulphur  compounds  in  paragraph  6e,  etc.  In  the  para- 
graph (9)  for  estimation  it  should  be  said,  nothing  more  than  a  general 
statement  of  methods  is  given,  for  the  benefit  of  qualitative  study,  with- 
out directions  and  specifications  for  quantitative  work,  in  which,  of 
course,  other  books  must  be  used. 

The  authors  desire  to  say  with  the  fullest  appreciation  that  Perry 
F.  Trowbridge,  instructor  in  Organic  Chemistry  in  this  University,  has 
performed  a  large  amount  of  labor  in  this  revision,  collecting  data  from 
original  authorities,  confirming  their  conclusions  by  his  own  experi- 
ments, elaborating  material,  and  making  researches  upon  questions  as 
they  have  arisen. 

University  of  Michigan, 
April,  1901. 


CONTENTS. 


PART  I.— THE  PRINCIPLES  OF  ANALYTICAL  CHEMISTRY. 

PAQH 

THE  CHEMICAL  ELEMENTS  AND   THEIR  ATOMIC  WEIGHTS 1 

TABLE  OF  THE  PERIODIC  SYSTEM  OP  THE  CHEMICAL  ELEMENTS 2 

DISCUSSION  OF  THE  PERIODIC  SYSTEM 3 

CLASSIFICATION  OF  THE   METALS  AS  BASEB 10 

COMMONLY  OCCURRING  ACIDS 13 

THE  OPERATIONS  OF  ANALYSIS 13 

SOLUTION  AND  IONIZATION 20 

ORDER  OF  LABORATORY  STUDY 25 

PART  II.— THE  METALS. 

THE  SILVER  AND  TIN  AND  COPPER  GROUPS. 

(FIRST  AND  SECOND  GROUPS). 

GENERAL  DISCUSSION 27 

THE  SILVER  GROUP  (FIRST  GROUP). 

Lead 29 

Mercury 37 

Silver 45 

Comparison  of  Certain   Reactions  of  the  Metals  of  the  Silver 

Group   " 51 

TABLE  FOR  ANALYSIS  OF  THE  SILVER  OR  FIRST  GROUP 52 

Directions  for  Analysis  with  Notes 53 

THE   TIN  AND   COPPER   GROUP  (SECOND   GROUP). 

THE    TIN    GROUP,  OR    SECOND    GROUP,  DIVISION   A. 

Arsenic 56 

Antimony 72 

Tin    82 

Comparison  of  Certain  Reactions  of  Arsenic,  Antimony  and  Tin.  90 

Gold 91 

Platinum 93 

Molybdenum 97 

THE    COPPER   GROUP,    OR   GROUP   II,  DIVISION   B. 

Bismuth 100 

Copper 104 

Cadmium ^ 110 

Comparison  of  Certain  Reactions  of  Bismuth,  Copper  and  Cad- 
mium    113 

Vii 


vin  CONTEXTS. 

PAGB 

THE  PRECIPITATION  OF  THE  METALS  OF  THE  SECOND  GROUP 113 

TABLE  FOR  THE  ANALYSIS  OF  THE  TIN  GROUP  (SECOND  GROUP,  DIVISION  A).  116 

Directions  for  Analysis  with  Notes 118 

TABLE  FOR  ANALYSIS  or  THE  COPPER  GROUP  (SECOND  GROUP,  DIVISION  B). .  126 
Directions  for  Analysis  with.   Notes -125 

RARER  METALS  OF  THE  TIN  AND  COPPER  GROUP. 

Ruthenium 131 

Rhodium 132 

Palladium 133 

Iridium 134 

Osmium 135 

Tungsten   136 

Germanium    137 

Tellurium    13g 

Selenium 139 

THE  IRON  AND  ZINC  GROUPS  (THIRD  AND  FOURTH  GROUPS) 141 

THE    IRON    GROUP    (THIRD    GROUP). 

Aluminum 144 

Chromium 148 

Iron 153 

TABLE  FOR  ANALYSIS  OF  THE  IRON  GROUP  (THIRD  GROUP) 163 

DIRECTIONS  FOR  ANALYSIS  WITH  NOTES 164 

THE  ZINC  GROUP  (FOURTH  GROUP). 

Cobalt   167 

Nickel 173 

Manganese 177 

Zinc 183 

Comparison  of  Some    Reactions  of  the   Iron    and    Zinc   Group 

Bases 187 

TABLE  FOR  THE  ANALYSIS  OF  THE  ZINC  GROUP  (FOURTH  GROUP) i8g 

DIRECTIONS  FOR  ANALYSIS  WITH  NOTES 189 

ANALYSIS  OF  IRON  AND  ZINC  GROUPS  AFTER  PRECIPITATION  BY  AMMONIUM 

SULPHIDE   191 

IRON  AND  ZINC  GROUPS  IN  PRESENCE  OF  PHOSPHATES 193 

IRON  AND  ZINC  GROUPS  IN  PRESENCE  OF  OXALATES 194 

Table  of  Separation  of  Iron,  Zinc  and  Calcium  Group  Metals 
and  Phosphoric  Acid  by  Means  of  Alkali  Acetate  and  Ferric 

Chloride 196 

Table  of  Separation  of  Iron,  Zinc  and  Calcium  Group  Metals 
and  Phosphoric  Acid  by  Means  of  Ferric  Chloride  and  Barium 
Carbonate 197 

THE  RARER  METALS  OF  THE  IRON  AND  ZINC  GROUPS. 

Cerium 198 

Columbium  (Niobium) 198 

Didymium 199 

Erbium 200 

Gallium..  .  200 


CONTENTS.  iX 

PAGE 

Glucinum  (Beryllium) 200 

Indium 201 

Lanthanum 202 

Neodymium 202 

Praseodymium 202 

Samarium 202 

Scandium 202 

Tantalum 203 

Terbium 203 

Thalliu in 204 

Thorium 204 

Titan  ium   205 

Uranium 206 

Vanadium 207 

Ytterbium 208 

Yttrium    208 

Zirconium 209 

THE  CALCIUM  GROUP  (FIFTH  GROUP).     (THE  ALKALINE  EARTH  METALS) 209 

Barium 211 

Strontium 214 

Calcium   216 

Magnesium 220 

TABLE  FOR  THE  ANALYSIS  OF  THE  CALCIUM  GROUP  (FIFTH  GROUP).  ......  223 

DIRECTION  TOR  ANALYSIS  WITH  NOTES 224 

SEPARATION  OF  BARIUM,  STRONTIUM,  AND  CALCIUM  BY  THE  USE  OF  ALCOHOL  226 

ALKALINE  EARTH  METALS  AS  PHOSPHATES 226 

ALKALINE  EARTH  METALS  AS  OXALATES 226 

THE  ALKALI  GROUP  (SIXTH  GROUP) 227 

Potassium 228 

Sodium 232 

Ammonium 235 

Caesium 239 

Rubidium 240 

Lithium   24') 

DIRECTIONS  FOR  ANALYSIS  WITH  NOTES 242 

PAET  III.— THE  NON-METALS. 

BALANCING  OF  EQUATIONS 246 

Hydrogen 250 

Boron 252 

Boric  Acid 252 

Carbon 254 

Acetic  Acid 256 

Citric  Acid 258 

Tartaric  Acid ." 259 

Carbon  Monoxide 262 

Oxalic  Acid 263 

Carbon  Dioxide  (Carbonates) 267 


CONTENTS. 

PAGE 

Cyanogen 271 

Hydrocyanic  Acid 271 

Hydroferrocyaiiic  Acid 275 

Hydroferricyanic  Acid 277 

Cyanic  Acid 279 

Thiocyanic  Acid 280 

Nitrogen 281 

Hydrazoic  Acid .282 

Nitrous  Oxide 283 

Nitric  Oxide 283 

Nitrous  Acid 284 

Nitrogen  Peroxide 285 

Nitric  Acid 285 

Oxygen   291 

Ozone 293 

Hydrogen  Peroxide 294 

Fluorine   297 

Hydrofluoric  Acid 298 

Fluosilicic  Acid 298 

Silicon 299 

Silicic  Acid 299 

Phosph  orus    301 

Phosphine 304 

Hypophosphorous  Acid 304 

Phosphorous  Acid 306 

Hypophosphoric  Acid  307 

Phosphoric  Acid 308 

Sulphur    313 

Hydrosulphuric  Acid 315 

Thiosulphuric  Acid 321 

Hyposulphurous  Acid 323 

Dithionic  Acid 324 

Trithionic  Acid 324 

Tetrathionic  Acid 325 

Pentathionic  Acid 325 

TABLE  OF  THIONIC  ACIDS 326 

Sulphurous  Acid 327 

Sulphuric  Acid 331 

Persulphuric  Acid 336 

Chlorine 337 

Hydrochloric  Acid 341 

Hypochlorous  Acid 348 

Chlorous   Acid 349 

Chlorine  Peroxide 353 

Chloric  Acid 350 

Perchloric  Acid 353 

Bromine 354 

Hydrobromic  Acid 357 

Hypobromous  Acid 360 


CONTENTS. 


XI 


PAGE 

Bromic  Acid 360 

Iodine   '. 362 

Hydriodic  Acid 365 

lodic  Acid 369 

Periodic  Acid 372 

COMPARATIVE  REACTIONS  OF  THE  HALOGEN  COMPOUNDS 373 

PART  IT.— SYSTEMATIC  EXAMINATIONS. 

REMOVAL  OF  ORGANIC   SUBSTANCES 374 

PRELIMINARY  EXAMINATION  OF  SOLIDS    375 

CONVERSION  OF  SOLIDS  INTO  LIQUIDS 378 

TREATMENT  OF  A  METAL  OR  AN  ALLOY  379 

SEPARATION  OF  ACIDS  FROM  BASES 381 

TABLE  FOR  PRELIMINARY  EXAMINATION  OF  SOLIDS 382 

BEHAVIOR  OF  SUBSTANCES  BEFORE  THE  BLOW-PIPE 386 

TABLE  OF  THE  GROUPING  OF  THE  METALS 387 

TABLE  FOR  THE  SEPARATION  OF  THE  METALS 388 

ACIDS— FIRST  TABLE 390 

ACIDS — SECOND  TABLE 398 

ACIDS— THIRD  TABLE t . .  399 

ACIDS— FOURTH  TABLE 400 

NOTES  ON  THE   DETECTION  OF  ACIDS 401 

PRINCIPLES 4Q5 

EQUATIONS 409 

PROBLEMS  IN  SYNTHESIS 410 

TABLE  OF  SOLUBILITIES 412 

REAGENTS 415 


ABBREVIATIONS. 


A. 

A.  Ch. 

Am. 

Am.  S. 

Arch.  Pharm. 
Am.  Chem. 

B. 

Bl. 

B.  J. 
Comey. 

C.  N. 
Ch.  Z. 
C.  r. 
C.  C. 
Dingl. 
D. 

Fehling. 
Fresenius. 
G.  O. 
Gazzetta. 
Gilb. 

Gmelin-Kraut. 
J. 

J.  C. 

J.  pr. 

J.  Soc.  Ind. 

J.  Anal. 

J.  Am.  Soc. 

J.  Pharm. 

Laden  burg. 

M. 

Phil.  Mag. 

Pogg. 

Proc.  Roy.  Soc. 

Pharm.   J.   Trans. 

Ph.  C. 

Tr. 

Watt's. 


1868* 


*  Indicates  continuance  to  the  present  time. 

Liebig's  Annalen.     1832* 

Annales  de  Chimie  et  de  Physique.     1789* 

American  Chemical  Journal.     1879* 

American  Journal  of  Science.     1818* 

Analyst.     1876* 

Archives  der  Pharmacie.     1822* 

American  Chemist.     1870-77. 

Berichte  der  Deutschen  Chemischen  Gesellschaft. 

Bulletin  de  la  Societe  Chimique.     1859* 

Berzelius  Jahresbericht.     1822-51. 

Comey's  Dictionary  of  Solubilities.     1896. 

Chemical  News.     1860* 

Chemiker  Zeitung.     1877* 

Comptes  Rendus  des  Seances  de  1' Academic  des  Sciences. 

Chemisches  Centralblatt.     1830*    - 

Dingler's  Polytechnische  Journal.     1820* 

Dammer's  Anorganische  Chemie.     1892* 

Fehling' s  Handbuch  der  Chemie.     1871* 

Fresenius:  Qualitative  Chemical  Analysis. 

Graham-Otto:  Lehrbuch  der  anorganischen  Chemie. 

Gazzetta  chimica  italiana.     1871* 

Gilbert's  Annalen  der  Physik  und  Chemie.     1799-1824. 

Gmelin-Kraut:  Handbuch  der  anorganischen  Chemie.     1877. 

Jahresbericht  iiber  die  Fortschritte  der  Chemie.     1847* 

Journal  of  the  Chemical  Society.     1849* 

Journal  fiir  praktische  Chemie.     1834* 

Journal  of  the  Society  of  Chemical  Industry.      1882* 

Journal  of  Analytical  Chemistry.     1887-1893. 

Journal  of  the  American  Chemical  Society.     1876* 

Journal  de  Pharmacie  et  de  Chimie.     1809» 

Handworterbuch  der  Chemie.     1882-1895. 

Monatshefte  fiir  Chemie.     1880* 

Menschutkin.     Locke's  Translation.     1895. 

Philosophical  Magazine.      1798* 

PoggendorfTs  Annalen  der  Physik  und  Chemie.     1824-1877. 

Proceedings  of  the  Royal  Society  of  London.     1832* 

Pharmaceutical  Journal  and  Transactions.     1841* 

Pharmaceutische  Centralhalle.     1859* 

Transactions  of  the  Royal  Society.     1665* 

Watt's  Dictionary  of  Chemistry.     1888. 


1835* 


Wells' Trans.,  1897. 
1885. 


COLLgat 


xiv  ABBREVIATIONS. 

W.  A.  Wiedemann's  Annalen.     1877* 

W.  A.  (Beibl.)  Wiedemann's  Annalen  Beiblatter.     1877* 

Wormley.  Wormley's  Microchemistry  of  Poisons.     1867. 

Wurtz.  Dictionnaire  de  Chimie.     1868. 

Z.  Zeitschrift  fiir  analytische  Chemie.     1862.* 

Z.  Ch.  Zeitschrift  fiir  Chemie.     1865-1871. 

Z.  anorg.  Zeitscbrift  fiir  anorganische  Chemie.     1891* 

Z.  angew.  Zeitscbrift  fiir  angewandte  Chemie.     1888* 

Z.  phys.  Ch.  Zeitschrift  fiir  physicaliscbe  Chemie.     1887* 


§1.    INTERNATIONAL  ATOMIC  WEIGHTS  FOR  1916. 

Compiled  by  the  International  Committee  on  Atomic 
Weights,  consisting  of  F.  W.  Clarke,  W.  Ostwald,  T.  E.  Thorpe, 
and  G.  Urbain, 


Aluminum 

Al 

O=16 

27  1 

Molybdenum  

O 

Mo 

=  .16. 

96  0 

Antimony    . 

Sb 

120.2 

Neodymium    

-Nd 

144  3 

Argon 

A 

39.88 

Neon  

Ne 

20  2 

Arsenic  

As 

74.96 

Nickel  

Ni 

58  6;5 

Barium    
Bismuth  
Boron 

Ba 
Bi 
B 

137.37 
208.0 
11  0 

Niton  (radium 
emanation)  
Nitrogen 

Nt 
N 

222.4 
14  01 

Bromine 

Br 

79  92 

Osmium. 

Os 

190  9 

Cadmium 

Cd 

112  40 

Oxygen 

o 

16  00 

CSBSI  im 

Cs 

132  81 

Palladium 

Pd 

106  7 

Calcium 

Ca 

40  07 

Phosphorus 

p 

31  04 

Carbon  

c 

12.00 

Platinum  

Pt 

195  2 

Cerium  
Chlorine  

Ce 
Cl 

140.25 
35.46 

Potassium  
Praseodymium  

K 
Pr 

39.10 
140  9 

Chrom1  um 

Cr 

5?  0 

Radium 

Ra 

226  0 

Cobalt 

Co 

58  97 

Rhodium 

Rh 

102  9 

Columbium 

Ch 

93  5 

Rubidium 

Rb 

85  45 

Copper 

Cn 

63  57 

Ruthenium 

Ru 

101  7 

Dysprosium  
Erbium  
Europium 

% 

Eu 

162.5 
167.7 
152  0 

Samarium  
Scandium  
Selenium               .    .  . 

Sa 
Sc 
Se 

150.4 
44.1 
79  2 

Fluorine 

F 

19  0 

Silicon                       .  . 

Si 

28.3 

Gadolinium. 

Gd 

157  3 

Silver 

Ag 

107  88 

Gallium    

Gn 

69  9 

Sodium                 

Nn 

23.00 

Germanium   

GP 

72  5 

Strontium          

Sr 

87.63 

Glucinum  

Gl 

9  1 

Sulfur.              

S 

32.06 

Gold  
Helium  

Au 
HP 

197.2 
4.00 

Tantalum  
Tellurium  

Ta 
Te 

181.5 
127.5 

Holmium  

Hrt 

163.5 

Terbium  

Tb 

159.2 

Hvdrogen  

H 

1.008 

Thallium  

71 

204.0; 

Indium 

In 

114  8 

Thorium 

Th 

232.4 

Iodine 

I 

126.92 

Thulium 

Tm 

168.5 

Iridium. 

Ir 

193.1 

Tin                        

Sn 

118.7 

Iron 

FP 

55.84 

Titanium.          

Ti 

48.1 

Krypton          

Kr 

82.92 

Tungsten  

W 

184.0 

Lanthanum.            .  . 

T* 

139.0 

Uranium      

u 

238.2 

Lead  

Pb 

207.20 

Vanadium  

V 

51.0 

Lithium  

Li 

6.94 

Xenon      

Xe 

130.2 

Lutecium 

Lu 

175  0 

Ytterbium  (Neoytter- 

Magnesium 

Me 

24  32 

bium)                    .  . 

Yb 

173  5 

Mangansse 

Mn 

54  93 

Yttrium                 .... 

Yt 

88  7 

Mercury  

Hr 

200  6 

Zinc                       .... 

Zn 

65.37 

Zirconium  

Zr 

90.6 

Jour.    Am.    Chem.    Soc.,    1915,    37,  2449. 


Th  00  t^ 

OOCOO5 


THOCO 


II  II  II 


O5rH 

:CO 
Oi  O5 


8  CO  ^O 


C^HH 


a 

H 
-Ji 

-si 

U 

q 

o 


so 


r-|  uj 


O 


0 


r-t  O 


PH  II 


10  g 

§      7 


o 


ii       °2 


O 


a 


o         S 


O 


0 


rH  CO 


O 


O 


CO 

8    °. 

(M        t& 


PS 


00 


OC 


00 


C5 


§3.  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  3 

§3.  In  the  periodic  system  of  the  chemical  elements  certain  regular 
gradations  of  chemical  character  are  to  be  studied  and  held  in  view,  to  sim- 
plify the  multitude  of  facts  observed  in  analysis.  Passing  from  Li  G.94  to 
F  19.0  in  the  first  Series  of  this  system,  the  elements  are  successively  less 
and  less  of  the  nature  to  constitute  bases  and  more  and  more  of  the  nature 
to  form  acids,  as  their  atomic  weights  increase.  The  acid-forming  elements 
are  electro-negative  to  the  elements  which  form  bases.* 

But  in  passing  from  F  19.0  to  the  next  higher  atomic  weight,  Na  23.0, 
we  return  from  the  acid  extreme  to  the  basal  extreme  and  begin  another 
period,  in  gradation  through  the  seven  Groups.  There  is  a  like  return 
from  one  extreme  to  the  other  in  the  steps  between  chlorine  and  potassium 

*  Bases  are  the  oxygen  compounds  of  the  metals.  Acids  are  compounds  of  elements  for 
the  most  part  not  metals.  In  the  chemical  union  of  sodium  with  chlorine,  for  example, 
these  two  elements  differ  widely  from  each  other  in  their  various  properties.  The  chlorine 
is  the  opposite  of  the  sodium  in  that  very  power  by  virtue  of  which  the  one  combines  with 
the  other  in  the  making  of  sodium  chloride,  a  distinct  product.  In  the  polarity  of  electro- 
lysis the  sodium  is  the  positive  element,  while  the  chlorine  is  the  negative  element.  The 
power  of  opposite  action  exercised  by  the  one  element  upon  the  other,  in  their  combination 
together,  is  represented  by  the  opposite  polarity  of  the  one  in  relation  to  the  other  during 
electrolysis.  Electrolysis  is  an  exercise  of  the  same  energy  that  is  otherwise  manifested 
in  chemical  union  or  in  a  chemical  change.  Strictly  speaking,  it  may  be  said  that  it  is  only 
in  electrical  results  that  a  positive  or  a  negative  polarity  appears.  But  the  term  positive 
polarity,  applied  to  sodium  because  it  goes  to  the  negative  pole  of  a  battery,  is  a  term 
which  well  designates  the  oppositeness  of  the  chemical  action  of  sodium  in  its  union  with 
chlorine.  That  is  to  say,  the  metals  are  in  general  "  positive,"  the  not-metals  in  general 
"  negative,"  in  the  relation  of  the  former  to  the  latter,  and  this  relation  may  be  termed 
one  of  "  polarity,"  whether  it  appear  in  electrolysis,  in  chemical  combination,  or  in  a 
chemical  change. 

In  chemical  combination,  the  atoms  of  each  element  act  with  a  "  polarity,"  the  extent 
of  which  may  be  expressed  in  terms  of  hydrogen  equivalence  or  "  valence."  The  valence  of 
an  element,  when  in  combination  with  another  element,  may  be  counted  as  relatively 
"  positive  "  or  "  negative  "  to  the  latter.  For  example,  in  the  compound  known  as  hydro- 
sulphuric  acid,  the  sulphur  is  negative,  the  hydrogen  positive,  in  the  relation  of  one  to  the 
other,  as  represented  by  the  diagram, 

H+- 

H+-S' 

in  which  the  plus  and  minus  signs  of  mathematics  are  used  to  represent  the  "  positive  " 
and  "  negative  "  activities  of  chemical  elements.  That  is,  the  sulphur  acts  with  two  units 
of.  valence,  both  in  negative  polarity.  In  sulphuric  acid  the  sulphur  is  positive  in  relation 
to  both  the  oxygen  and  the  hydroxyl,  as  indicated  in  the  diagram 

(H0)-+        I   +-° 


That  is,  the  sulphur  acts  with  six  units  of  valence,  all  in  positive  polarity.  In  respect  to 
oxidation  and  reduction,  the  difference  between  the  action  of  sulphur  in  hydrosulphuric 
acid  on  the  one  hand,  and  in  sulphuric  acid  on  the  other  hand,  is  a  difference  equivalent  to 
eight  units  of  -valence,  the  combining  extent  of  eight  atoms  of  hydrogen.  This  value  is  in 
agreement  with  the  factors  of  oxidizing  agents  in  volumetric  analysis. 

In  the  same  sense  there  is  a  change  of  "  polarity  "  equivalent  to  the  extent  of  eight  units 
of  valence,  in  reducing  periodic  acid  to  hydriodic  acid,  in  reducing  arsenic  acid  to  arsine,  or 
in  reducing  carbon  tetrachloride  to  methane.  That  is,  in  any  of  the  groups  from  IV.  to 
VII.  there  is  a  difference,  equivalent  to  the  combining  extent  of  eight  hydrogen  units,  be- 
tween the  negative  polarity  of  the  element  in  its  regular  combination  with  hydrogen,  such 
as  NH?,  and  its  positive  polarity  in  its  highest  combination  with  oxygen,  such  as  NOa  (OH). 


4  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  §4. 

and  in  those  between  bromine  and  rubidium.  This  fact  of  a  periodic 
return  in  the  gradation  of  the  properties  of  the  elements,  as  their  atomic 
weights  ascend,  constitutes  a  periodic  system.  A  period  is  termed  a  Series. 
A  Group  in  this  system  consists  of  the  corresponding  members  of  all  the 
Series,  which  members  are  found  to  agree  in  valence,  so  that  the  number 
of  the  groups,  from  I.  to  VII.  (not  in  VIII.),  expresses  the  typical 
valence  of  the  elements  as  grouped.  Further  inquiry  shows  that  all  the 
properties  of  the  elements  are  in  relation  to  their  atomic  weights,  as  they 
appear  in  the  periodic  system.  But  this  system  is  not  to  be  depended  upon 
to  give  information  of  the  facts;  it  is  rather  to  be  used  as  a  compact  simpli- 
fication of  facts  found  independently,  by  the  student  and  by  the  author- 
ities on  whom  the  student  must  depend.  A  full  account  of  the  Periodic 
System,  as  far  as  it  is  understood,  is  left  to  works  on  General  Chemistry. 

§4.  The  remarkable  position  of  Group  VIII.,  made  up  of  three  series. 
each  of  three  elements  near  each  other  in  atomic  weight,  respectively  in 
Series  4,  6,  and  10,  is  in  central  relation  to  the  entire  system.  In  this 
group  there  is  something  of  a  return,  from  negative  to  positive  polarity, 
from  higher  to  lower  valence.  Group  VIII.  lies  between  Group  VII.  and 
Group  I.,  that  is  to  say  in  this  group  there  is  a  return  from  negative  to 
positive  nature,  and  from  higher  to  lower  valence.  Moreover,  the  newly 
discovered  elements  related  to  argon,  destitute  of  combining  value  as  they 
are,  appear  to  constitute  a  Group  0.  The  latest  results  render  this  position 
of  the  argon  group  of  elements  so  probable  that  it  has  been  placed  in  the 
chart  for  convenience  of  study,  subject  to  further  conclusions.  (W.  Ramsay. 
Br.  Assoc.  Adv.  Sci.,  1897,  598-601;  B.  1898,  31,  3111.  J.  L.  Howe,  C.  N.9 
1899,  80,  74;  1900,  82,  15,  52.  Ostwald,  Grundr.  Allg.  Chem.,  3te  Auf., 
1899,  S.  45.)  In  comparison  with  the  members  of  Group  VII.  those  of 
Group  VIII.  certainly  have  a  diminished  negative  polarity,  and  a  lower 
valence,  the  latter  being  easily  variable.  Some  of  the  particulars  are  given 
below  under  the  head,  "  Metals  in  Relation  to  Iron."  The  most  remark- 
able thing  about  Group  VIII.  is  the  fact  that  the  return  to  Group  I.  from 
Group  VIII.  is  less  complete  than  the  return  from  Group  VII.  That  is  to 
say,  the  character  of  copper  is  divided  between  Group  VIII.  and  Group  I.. 
and  the  same  is  true  of  silver  and  of  gold.  This  relation  to  Group  VIII. 
can  be  traced,  in  some  particulars,  to  zinc  and  cadmium  and  mercury  in 
Group  II.  For  these  reasons  Series  4  and  5  may  be  studied  as  one  long 
period  of  seventeen  members,  Series  6  and  7  as  another  long  period  and 
Series  10  and  11  as  a  third  and  final  long  period. 

§5.  It  is  to  be  observed  that  each  one  of  the  Groups,  from  I.  to  VII.,  falJa 
in  two  columns,  a  column  consisting  of  the  alternate  elements  in  the  group. 
Thus,  H,  Li,  K,  Rb  and  Cs  make  up  the  first  column  of  Group  I.  It  is 
among  the  alternate  members  of  a  group  that  the  closer  grade-relations  of 


§9.  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  5 

the  elements  are  found.  The  gradations  represented  under  one  column 
are  distinct  from  those  under  the  other  in  the  same  group.  The  well 
known  alternate  elements  of  a  Group,  so  far  as  found  clearly  graded 
together  in  respect  to  given  properties,  are  to  be  studied  as  a  Family  of 
elements.  Again  a  number  of  elements  next  each  other  in  a  Series  are  to 
be  studied  together,  either  by  themselves  or  with  an  adjoining  half-group. 

For  the  studies  of  analytical  chemistry  the  following  are  the  more 
strongly  marked  of  the  families  of  the  well  known  elements. 

§6.  The  Alkali  Metals.— 116.94,  (Na  23.0),  K  39.10,  Rb  85.45,  Csl32.81. 
The  first  part  and  sodium  of  the  second  part  of  Group  I.  In  the  grada- 
tion of  these  elements  the  basal  power  increases  qualitatively  with  the  rise 
in  atomic  weight.  The  hydroxides  and  nearly  all  salts  of  these  metals  are 
freely  soluble  in  water,  wherein  they  are  unlike  the  ordinary  metals  of  all 
the  other  groups.  For  the  most  part,  however,  these  solubilities  increase 
with  the  atomic  weight  of  the  metal,  and  the  carbonate  and  orthophosphate 
of  lithium  are  but  slightly  soluble. 

§7.  The  Alkaline  Earth  Metals.— (Mg  24.32),  Ca  40.07,  Sr87.63;Ba  137.37. 
These  metals,  like  those  of  the  alkalis,  form  stronger  bases  as  they  have 
higher  atomic  weights.  Both  in  Group  I.  and  in  Group  II.  the  member 
in  Series  3  (Na,  Mg),  though  in  the  second  set  of  alternate  members,  agrees 
in  many  ways  with  the  next  three  of  the  first  set  of  alternates.  The 
hydroxides  of  these  metals  are  not  freely  soluble  in  water  but  arq  regularly 
more  soluble  as  the  atomic  weight  of  the  metal  is  higher.  The  sulphides 
are  freely  soluble;  the  carbonates  and  orthophosphates  quite  insoluble. 
The  sulphates  have  a  graded  solubility,  decreasing  as  the  atomic  weight 
is  higher,  an  order  of  gradation  the  reverse  of  that  of  the  hydroxides  and 
of  wider  range.  That  is,  at  one  extreme  the  magnesium  sulphate  is  freely 
soluble,  at  the  other  barium  sulphate  is  insoluble. 

§8.   The  Zinc  Family.—  Mg  24.32,  (Al  27.1),  Zn  65.37,  Cd  112.4, , 

H<*  200.6.  These  metals,  save  aluminum,  belong  to  the  second  alternates  of 
Group  II.,  and,  like  those  of  the  corresponding  half  of  Group  I.,  in  their 
gradation  they  are  in  general  less  strongly  basal  as  they  rise  in  their  atomic 
weights.  Aluminum,  here  drawn  in  from  Group  III.  second  half,  has  the 
valence  of  the  third  group,  and  differs  from  the  others  in  not  forming  a 
sulphide.  The  sulphide  of  magnesium  is  soluble,  the  sulphides  of  zinc, 
cadmium  and  mercury  insoluble  in  water,  and  these  three  show  this  grada- 
tion, that  the  zinc  sulphide  is  the  one  dissolved  by  dilute  acid,  while  the 
mercury  sulphide  is  the  one  requiring  a  special  strong  acid  to  dissolve  it. 
both  these  differences  being  depended  upon  in  analysis.  Mercury,  sepa- 
rated from  cadmium  by  two  removes  in  the  periodic  order,  is  but  a  distant 
member  of  this  family. 

§9.  Metals  in  Relation  to  Iron.—Cr  52.0,  Mn  54.93,  Fe  55.84,  Ni  58.68, 


6  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  §10, 

Co  58.97.    The  atomic  weights  of  these  metals  lie  nearly  together.    They 
all  belong  to  one  Series,  the  fourth,  representing  Groups  VI.  and  VII., 
and  make  the  first  of  the  instances  of  three  members  together  in  one  seriei 
in  Group  VIII.     Chromium,  being  in  the  first  division  of  its  group,  could 
not  be  expected  to  grade  with  sulphur  and  selenium,  nor  would  manganese 
be  expected  to  grade  with  chlorine  and  bromine,  but  the  disparity  is  strik- 
ing in  both  cases,  especially  in  the  comparison  of  melting  points.     The 
valence  of  both  chromium  and  manganese  appears  partly  exceptional  to 
their  positions  in  the  system  but  the  maximum  valence  of  each  is  regular, 
That  all  of  these  five  elements,  neighbors  to  chlorine  and  bromine,  are 
counted  as  metals,  is  not  contrary  to  the  periodic  order.     Group  VIII.  binds 
Group  I.  to  Group  VII.     After  Co  58.97  follow  Cu  63.57  and  then  Zn  65.37. 
Indeed  each  of  "  the  well-known  metals  related  to  iron  "  is  capable  of  serv- 
ing as  either  a  base  or  an  acid,  by  change  of  valence.     These  metals  are  the 
special  subjects  of  oxidation  and  reduction.     So  far  they  resemble  their 
non-metallic  neighbors,  the  halogens.     Of  the  five,  chromium  and  man- 
ganese (nearest  the  halogens)   form  the  best  known  acids.     Nickel  and 
cobalt,  like  copper,  have  a  narrower  range  of  valence,  a  more  limited  extent 
of   oxidation   and   reduction,   within   which   they  as   readily   act.     These 
valences,  in  capacity  of  combination  with  other  elements,  not  including  the 
most  unusual  valences,  may  be  written  in  symbols  as  follows: 

2-3-6  2-3-4-6-7  2-3-6  2-3  2-3  1-2  2 

Cr     ,  Mn     ,  JFe     ,  Ni     ,  Co     ,  Cu     ,  Zn 

On  reaching  zinc,  65.37,  in  this  gradation,  the  capacity  of  oxidation  and 
reduction  disappears.  Sulphides  are  formed  by  such  of  these  metals  as  act 
with  a  valence  of  two  (all  except  chromium),  and  these  sulphides  are  insolu- 
ble in  water.  In  the  conditions  of  precipitation  sulphides  are  not  formed 
with  the  metal  in  any  valence  other  than  two.  Chromium  acting  as  a 
base  with  a  valence  of  three,  like  aluminum  whose  only  valence  is  three, 
refuses  to  unite  with  sulphur.  Trivalent  iron  in  precipitation  by  sulphides 
is  mainly  reduced  to  ferrous  sulphide  (FeS).  In  chromates  the  chromium 
valence  is  reduced  from  six  to  three  by  hydrogen  sulphide  acting  in  solu- 
tion. A  carbonate  is  not  formed  by  chromium,  this  being  another  agree- 
ment with  aluminum,  and  the  same  is  true  of  trivalent  iron.* 

§10.  The  Metals  not  Alkalis  in  Group  I.,  Second  Part,  and  their  Relatives 

in  Group  VIII. — Cu  63.57,  Ag  107.88, ,  Au  197.2.  In  gradation  these 

metals  are  less  strongly  basal,  and  more  easily  reduced  from  their  com- 
pounds to  the  metallic  state,  as  their  atomic  weights  rise.  This  is  in  agree- 
ment with  the  gradation  among  the  second  set  of  alternates  in  Group  II.. 
the  Zinc  Family.  It  likewise  agrees  with  second  part  of  Group  VII. ,  the 
halogens.  These  elements  of  Group  I.  are  to  be  studied  with  those  of 
Group  VIII.,  especially  with  those  respectively  nearest  them  in  atomic 

*  These  metals  form  unstable  hydrated  basic  carbonates. 


§12.  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  7 

weight:  Cu  63.57  with  Ni  58.68  and  Co  58.97,  Ag  107.88  with  Pd  106. 7,  and 
Au  197.2  with  Pt  195.2.  Those  with  atomic  weights  above  that  of  copper 
rank  as  "  noble  metals/'  from  their  resistance  to  oxidation  and  other 
qualities,  so  ranking  in  higher  degree  as  their  atomic  weights  increase. 
Their  melting  points  (those  of  Pd,  Ag,  Au,  Pt)  rise  in  the  same  gradation. 

By  action  of  ammonium  hydroxide  upon  solutions  of  their  salts  these 
(seven)  metals  form  metal  ammonium  compounds,  all  of  which  are  soluble 
in  water  except  the  compounds  of  platinum  and  gold  (highest  in  atomic 
weight).  All  of  the  seven  named  form  sulphides  insoluble  in  water,  in 
condition  of  precipitation.  For  the  most  part  their  sulphides  are  relatively 
more  stable  than  their  oxides.  Silver  differs  from  the  others  in  the  insolu- 
bility of  its  chloride,  and  agrees  irregularly  in  this  fact,  one  prominent  in 
analysis,  with  mercury  in  its  lower  valence,  and  partly  with  lead. 

§11.   The  Nitrogen  Family  of  Elements.— -N  14.01,  P  31.04,  As  74.96, 

Sb  120.2,  ,  Bi  208.0.     These  elements  include  the  leading  element  of 

Group  V.,  and  the  entire  second  part  of  the  group.  Nitrogen  and  phos- 
phorus act  as  non-metals,  antimony  and  bismuth  as  metals,  while  arsenic  is 
intermediate,  the  polarity  being  more  positive  as  the  atomic  weight  increases. 
In  combinations  with  hydrogen,  like  ammonia  and  ammonium  compounds, 
phosphine  and  phosphonium  salts,  and  also  like  analogous  organic  bases 
where  carbo-hydrogen  takes  the  place  of  a  part  or  all  of  the  hydrogen,  there 
is  a  remarkable  unity  of  type  in  this  family.  The  same  is  true  of  the  com- 
binations with  oxygen,  like  nitric  acid.  It  is  in  Group  V.  that  the  group 
valence  for  oxygen  begins  to  diverge  in  gradation  from  the  group  valence 
for  hydrogen.  While  in  ammonium  compounds  nitrogen  exercises  a  valence 
of  five,  this  total  of  five  units  is 'always  limited  in  polarity  to  a  balance  of 
three  negative  units  at  most.  In  ammonia:  N~3.HHH.  In  ammonium 
chloride:  N~4+1=~3 .  HHHHG1.  Bismuth  is  a  distant  member,  a  vacancy 
falling  between  it  and  antimony. 

Phosphorus,  arsenic  and  antimony  are  in  gradation  with  each  other  as  to 
their  indifference  to  chemical  combination  and  readiness  of  reduction  to  the 
elemental  state,  these  qualities  intensifying  with  the  rise  in  atomic  weight. 
In  this  gradation  nitrogen,  belonging  among  the  other  alternate  members, 
has  no  part.  In  its  chemical  indifference  it  stands  in  extreme  contrast  to 
phosphorus. 

§12.  'Relation  of  Tin  and  Lead  to  the  Nitrogen  Family. — These  metals 
are  in  Group  IV.,  having  valences  of  four  and  five,  differing  from  the  valence 
of  the  nitrogen  family.  In  Series  7:  Sn  118.7  is  closely  related  to  Sb  120.2. 
In  Series  11:  Pb  207.20  is  closely  related  to  Bi  208.0.  The  metals  in  the 
first  named  pair  are  two  removes  from  those  in  the  second  pair,  all  being 
among  the  second  alternate  members.  In  their  salts  tin  and  antimony  are 
more  easily  subject  to  changes  of  valence  than  are  lead  and  bismuth.  In 

riti  irnDMiA    ftftl  1  EfiE 


8  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  §  12. 

further  comarison,  arsenic,  in  its  deportment  as  a  metal,  may  be  included 
making  the  list:  As  74.96,  Sb  120.2  (Sn  118.7),  Bi  208.0  (Pb  207.20).  Of 
these,  only  arsenic  forms  a  higher  oxide  soluble  in  water  (separation  after 
treatment  with  nitric  acid  and  evaporation).  Arsenic  and  antimony  form 
gaseous  hydrides,  in  this  agreeing  with  phosphorus  and  nitrogen,  while  the 
others  do  not.  The  stability  of  the  hydrides  of  N,  P,  As,  Sb,  all  in  the  type 
of  ammonia,  is  in  the  ratio  inverse  to  that  of  the  atomic  weight.  All  of 
these  metals  (As,  Sb,  Sn,  Bi,  Pb)  are  precipitable  as  hydroxides  save  arsenic, 
all  are  precipitated  as  sulphides,  and  these  have  chemical  solubilities  some- 
what in  gradation  with  atomic  weights,  the  arsenic  sulphide  being  most 
fully  separable  by  chemical  solvents.  The  sparing  solubility  of  the  chloride 
of  lead,  referred  to  in  description  of  silver,  is  approached  by  the  insolubility 
of  the  oxychlorides  of  bismuth,  tin,  and  antimony,  and  this  fact  must  be 
borne  in  mind,  when  precipitation  by  hydrochloric  acid  is  employed  for 
the  separation  of  silver  and  univalent  mercury  in  analysis. 

Nitrogen  in  its  trivalent  union  with  hydrogen,  the  leading  element  of  the 
group  of  alkali  metals,  constitutes  an  active  alkali.  In  its  prevalent  union  with 
oxygen,  the  leading  element  of  Group  VI. ,  that  is  with  oxygen  and  hydroxyl, 
nitrogen  forms  an  acid  which  is  very  active  though  not  very  stable,  its  decom- 
position being  represented  by  gunpowder.  The  degree  of  negative  polarity  of 
nitrogen,  or  its  capacity  for  acid  formation,  in  accordance  with  its  place  next 
to  oxygen  among  the  atomic  weights,  is  shown  in  that  singular  unstable  body, 
hydrazoic  acid,  HN3  (also  called  azimide),  of  decided  acid  power,  constitut- 
ing well  marked  salts,  such  as  sodium  azoimide,  Na  N3 ,  in  which  a  ring 
of  nitrogen  alone  acts  as  an  acid  radical.  The  first  four  members  of  the 
nitrogen  family  agree  with  each  other  in  forming  trivalent  and  pentavalent 
anhydrides  and  acids,  the  pentavalent  ones  being  the  more  stable.  The 
pentavalent  acids  are  of  especial  interest.  In  nitric  acid  the  five  units  of 
positive  valence  of  an  atom  of  nitrogen  are  met  by  two  atoms  of  oxygen 
with  two  units  each  of  negative  valence  and  a  unit  of  negative  valence 

of  hydroxyl:  H — 0 — NZlQ .     The  same  constitution  is  found  in  metaphos- 

phoric  acid  HO  P  02  ,  meta-arsenic  acid  HO  As  02 ,  and  in  antimonic  acid 
HO  Sb  02.  The  so-called  orthq  acids,  phosphoric  and  arsenic,  have  the 
constitution  (H0)3  P  0  and  (H0)3  As  0  ,  respectively.  Phosphoric  and 
arsenic  acids  have  a  remarkable  likeness  to  each  other  in  nearly  all  the 
properties  of  all  their  salts,  behaving  alike  in  analysis  so  long  as  preserved 
from  the  action  of  reducing  agents.  These  sharply  separate  arsenic,  usually  in 
one  of  its  trivalent  forms,  AsH3  or  As2S3 .  Antimony  is  reduced  from  its 
acid  even  more  readily  than  is  arsenic,  in  accordance  with  the  gradation 
stated  above. 

In  the  solubility  of  its  salts  with  metals,  the  acid  of  nitrogen  is,  again,  in 


§14.  DISCUSSION  OF  THE  PERIODIC  SYSTEM.  9 

strong  contrast  with  the  acids  of  the  elements  of  the  second  part,  phos- 
phoric and  arsenic  acids.  Metal  nitrates  are  generally  all  soluble  in  water. 
Of  the  metal  phosphates  and  arsenates,  that  is  the  full  metallic  salts  of 
phosphoric  and  arsenic  acids,  in  their  several  forms,  only  those  of  the  alkali 
metals  dissolve  in  water. 

§13.  TJie  Halogens.—*  19.00,  Cl  35.46,  Br  79.92,  I  120  92.'  These 
metals  constitute  the  leading  elements  of  Group  VI f.,  and  the  three 
known  members  of  the  second  alternates.  In  the  halogen  family 
fluorine  has  a  relation  like  that  of  nitrogen  in  its  family,  taking  part 
in  trie  group  gradation  as  to  polarity,  solubility  of  compounds  and 
other  qualities,  but  standing  quite  by  itself  in  respect  to  certain  properties. 
It  is  the  most  strongly  electro-negative  of  the  known  elements,  a  fact  in 
accord  with  the  relation  of  its  atomic  weight. 

For  the  common  work  of  analysis  we  may  confine  our  study  of  the 
halogens  to  chlorine,  bromine,  and  iodine.  In  the  order  of  their  atomic 
weights,  these  elements  appear,  respectively,  in  gaseous,  liquid,  and  solid 
state,  under  common  conditions.  Their  hydrogen  acids,  HC1 ,  HBr ,  and 
HI,  show  a  stability  in  proportion  to  the  electro-negative  polarity  of  the 
halogen,  hydriodic  acid  being  so  unstable  as  to  suffer  decomposition  in  the 
air.  In  the  solubility  of  their  metal  salts  these  acids  are  nearly  alike,  all 
being  soluble  except  the  silver,  univalent  mercury,  and  lead  salts,  but  the 
iodides  of  divalent  mercury,  bismuth  and  divalent  palladium  are  sparingly 
soluble.  Each  of  these  halogens,  most  especially  iodine,  forms  a  class  of 
salts  each  containing  two  metals,  one  of  the  united  metals  being  that  of  an 
alkali,  such  as  (KI)2  HgI2  and  K2  Pt  C16 .  The  periodides  show  that  iodine 
atoms  have  the  power  of  uniting  with  each  other,  in  the  molecules  of  salts, 
a  power  partly  shared  by  bromine  and  chlorine  and  probably  exercised  in 
many  complex  halogen  compounds.  By  this  means  two  atoms  of  a  halogen 
may  serve  the  same  as  one  atom  of  oxygen,  in  the  linkings  of  molecular 
structure. 

Of  the  oxygen  acids  of  chlorine,  bromine  and- iodine,  those  in  which  the 
halogen  has  a  valence  of  five  are  more  stable  than  the  others.  These  acids 
are  chloric,  HO  Cl  02 ;  bromic,  HO  Br  02 ;  and  iodic,  HO  I  02.  Chloric  acid 
resembles  nitric  acid,  HO  N  02 ,  in  the  fact  that  it  forms  soluble  salts  with 
all  the  metals.  Chlorates  decompose  more  violently  than  nitrates;  iodates 
for  the  most  part  less  readily  than  the  latter.  Of  the  oxygon  acids  with 
a  halogen  valence  of  seven,  periodic  acid,  HO  I  03 ,  also  (H0)5 1  0  ,  is  pre- 
served intact  without  difficulty.  Perchloric  acid  is  more  stable  than  chloric  acid. 

§14.  The  Relations  of  Sulphur.— $  32.07.  Sulphur  is  the  first  member 
of  a  family  including  selenium  and  tellurium.  It  differs  from  oxygen 
almost  as  much  as  phosphorus  differs  from  nitrogen,  and  we  may  say  more 
than  silicon  differs  from  carbon.  The  higher  valence  of  Group  VI.,  exer- 


10  THE  CLASSIFICATION  OF  THE  METALS  AS  BASES.  §15. 

cised  toward  oxygen,  cannot  be  met  by  oxygen  itself.  Of  the  acids  of 
sulphur,  H2S ,  in  which  sulphur  has  two  electro-negative  units  of  valence, 
is  quite  unstable,  while  (H0)2  S  02 ,  in  which  the  sulphur  has  six  electro- 
positive units  of  valence,  is  the  most  stable.  The  sulphides  (salts  of  H2S) 
of  the  heavier  metals  quite  generally  are  insoluble  in  water,  an  important 
means  of  separation  in  analysis.  The  sulphates  (salts  of  H2S04)  of  the 
larger  number  of  the  metals  are  soluble  in  water,  the  exceptions  being 
important  to  observe,  namely  those  of  Pb  207.20, Ba  i:>7.:>7,  Sr  87.63,  and 
(with  sparing  solubility)  Ca  40.07.  Of  these  sulphates,  that  of  barium 
(least  soluble),  is  the  one  usually  employed  in  analytical  separation. 

§15.  The  Relations  of  Carbon. — C  12.0.  Carbon,  in  a  central  position 
in  respect  to  polarity,  stands  alone  in  its  capacity  for  a  multitude  of  dis- 
tinct compounds  with  hydrogen  and  oxygen,  with  and  without  nitrogen, 
these  being  the  so-called  organic  compounds.  This  capacity  goes  with 
the  power  of  carbon  atoms  to  unite  with  each  other  in  the  same  mole- 
cule. It  appears  in  acetylene  C2  H2  (H  C  EEC  H),  also  in  oxalic  acid, 
(HO)  OC  —  CO  (OH),  The  same  capacity  of  union  of  the  atoms  of  an 
element  with  each  other,  in  the  molecules  of  compounds,  is  exercised 
by  other  elements  in  fewer  instances,  as  by  nitrogen  in  hydrazoic  acid,  by 
oxygen  in  ozone,  by  sulphur  in  thiosulphuric  acid,  and  by  iodine  in  perio- 
dides.  In  carbon,  nitrogen,  and  oxygen  we  see  a  decreasing  gradation  01 
this  capacity,  as  the  atomic  weights  ascend.  Silicon,  next  to  caroon  in 
Group  IV.,  but  in  the  opposite  set  of  alternates,  agrees  with  carbon  m  tne 
formation  of  many  corresponding  compounds,  and  also  exhibits  to  some 
extent  the  capacity  of  uniting  its  atoms  to  each  otuer  in  ounding  up  com- 
binations. 


§16.  The  Classification  of  the  Metals  as  liases. 

The  object  of  the  Periodic  System  is  to  group  .ill  the  elements,  both 
metallic  and  non-metallic,  according  to  their  general  properties  as  related 
to  their  atomic  weights.  This  has  been  briefly  given  in  the  foregoing 
pages  for  study  bearing  especially  upon  the  main  methods  of  analysis. 

The  ordinary  grouping  of  the  bases  in  the  work  of  analysis,  outlined  in 
the  next  paragraph,  is  done  by  the  action  of  a  few  chemical  agents,  termed 
"group  reagents,"  which  have  been  chosen  from  a  large  number  of  re- 
agents, as  being  more  satisfactory  than  others,  for  the  use  of  the  greater 
number  of  analysts.  This  ordinary  grouping,  therefore,  is  not  the  only 
way  in  which  the  metals  can  be  separated,  in  the  practice  of  analytical 
chemistry,  nor  is  any  one  scheme  of  separation  adopted  throughout  by  all 
authorities.  The  principal  separations  of  analysis  can  be  well  understood 
by  gaining  an  acquaintance  witli  the  properties  of  the  leading  bases  and  acids. 


JJ16. 


THE  CLASSIFICATION  OF  THE  METALS  AS  BASES. 


11 


in  their  action  upon  each  other.     Without  this  acquaintance,  the  analyst  is 
the  servant  of  routine,  and  his  results  liable  to  fallacy. 

The  following  named  are  the  bases  of  more  common  occurrence. 


Metals  Precipitated  as  Chlorides, 
ine  Silver  Group. 
The  first  group.* 

Silver  (Argentum). 
f  Ag1:  silver  salts. 

Mercury  (Hydrargyrum) , 
Hff1:  merer  rous  salts. 

Lead  (Plumbum). 
Ph11;  lead  salts, 


Silver  and  the  mercury  of  mereur- 
ons  salts  can  be  removed  as  chlorides 
by  precipitation  with  hydrochloric 
acid.  The  precipitate  of  lead  is  not 
insoluble  enough  to  remove  this 
metal  entirely  in  separation  from 
other  groups. 


Metals  falling  with  Copper  and  Tin.         Precipitated  ~by  H.^S  in  acidulated 

The  second  group.  solution.     (The  precipitates  are  sul- 

phides.) 


The  Tin  Group. 

Division  A,  second  grout). 
bnxi  •  stannous  salts. 
SniV :  stannic  salts  and  stannates. 
Sbm:  antimonous  compounds. 
Sbv:  antimonic  compounds. 
As111:  arsenous  compounds. 
Asv:  arsenic  compounds  and  arsen- 
ates. 


Separated  by  dissolving  the  precip- 
itated sulphides  with  Ammonium 
Sulphide. 


Separated  by  the  insolubility  of  the 
precipitated  sulphides  on  treatment 
with  Ammonium  Sulphide. 


The  Copper  Group. 

Division  B,  second  groun. 
Hgn :  mercuric  salts. 
Pb11:  lead  salts. 
Bim :  bismuth  salts. 
Cu11:  copper  or  cupric  salts. 
Cu1:  cuprous  salts. 
Cd11 :  cadmium  salts. 


*  The  first  division  of  the  bases,  in  the  ordar  in  which  they  are  separated  from  each  other  by 
precipitation  with  the  group  reagents. 

t  The  Roman  numerals  (as  J)  express  units  o*  valence,  each  equivalent  to  an  atom  of  hydrogen, 
in  the  formation  of  salts  and  other  combinations. 


CLASSIFICATION  OF  THE  METALS  AS  BASES. 


§16. 


The  Iron  Group. 

The  third  group. 
Fe11  :  ferrous  salts. 
Fem:  ferric  salts. 
Crm:  chromic  salts. 
Crvl:  chrornates. 
Alm:  aluminium  salts. 

The  Zinc  Group. 

The  fourth  group. 
Zn11  :  zinc  salts. 
MnIT  :  manganous  salts. 

manganic  salts. 

unstable  salts. 

salts  of  manganic  acid. 

salts  of  permanganic  acid 
Ni11:  nickel  salts. 
Co11:  cobaltous  salts. 
Co111  :  cobaltic  salts. 

The  Alkaline  Earth  Bases, 
The  fifth  group. 

Magnesium,  Mg11. 


Mn111 
MnIV 


Calcium,  Calr. 
Strontium,  Srn. 
Barium,  Ban. 


The  Alkali  Bases. 

Tlie  sixth  group. 


Separated  by  precipitation  ivith 
Ammonium  Hydroxide,  in  presence 
of  NH,C1,  after  the  removal  of  previ- 
ously named  groups.  (The  precip- 
itates are  all  hydroxides.) 


Separated  by  precipitation  with 
Ammonium  Sulphide,  after  remoi.u 
of  all  previously  named  oases,  as  di- 
rected above.  (The  precipuates  are 
all  sulphides.) 


(Precipitated  by  carbonates,  wh'ch 
fact  alone  does  not  separate  tnem 
from  the  following  named  groups.) 

Separated  by  precipitation  as  ti 
phosphate  after  removing  all  the  pre- 
viously named  bases.  Forms  magne- 
sium hydroxide,  Mg(OH),,  and  mag- 
nesium salts,  such  as  MgS04 . 

Separated  by  precipitation  with 
Ammonium  Carbonate,  adding  NH^Cl 
to  keep  magnesium  from  precipitation. 
Calcium  carbonate,  a  normal  salt, 
CaCO  3. 

Not  precipitated  from  their  salts  by 
any  of  the  group  reagents.  Potassium 
and  sodium  are  found  after  removing 
aa  the  previously  named  groups. 
Ammonium  is  found  by  tests  of  the 
original,  this  base  being  added  in  the 
"group  reagents." 


§18. 

Potassium  (Kaliurn),  K1. 


THE  OPERATIONS  OF  ANALYSIS. 


13 


In  combination  in  potassium  hy- 
droxide, KOH,  and  in  potassium  salts, 
such  as  the  chloride  KC1,  and  the  ni- 
trate KNO, . 


Sodium  (Natrium), 


In  the  base,  sodium  hydroxide  and 
its  salts. 


Ammonium 


Forms  ammonium  hydroxide^ 
NH4OH,  representing  ammonia,  NHj, 
and  water,  and  serving  as  the  base  of 
ammonium  salts,  such  as 
ammonium  sulphate. 


§17.  THE  ACIDS  or  CERTAIN  COMMONLY  OCCURRING  SALTS. 


Name  of  Acid. 

Name  of  Salt. 

Formula. 

Showing  Hydroxyl. 

Anhydride. 

Carbonic 

Carbonate 

H2C03 

(HO)2CivO 

C02 

Oxalic 

Oxalate 

H2C204 

(HO)2C2iv02 

C203 

Nitric 

Nitrate 

HN03 

(HO)NV02 

N205 

Nitrous 

Nitrite 

HN02 

(HO)NniO 

N203 

Phosphoric  (ortho) 

Phosphate 

H3P04 

(HO)3PvO 

P205 

Metaphosphoric 

Metaphosphate 

HP03 

(HO)Pv02 

P20s 

Pyrophosphoric 

Pyrophosphate 

H4P20T 

(HO)4PV203 

P205 

Sulphuric 

Sulphate 

H2S04 

(HO)2Svi"02 

S03 

Sulphurous 

Sulphite 

H2S03 

S02 

Hydrosulphuric 

Sulphide 

H2S 

Hydrochloric 

Chloride 

HC1 

Hydrobromic 

Bromide 

HBr 

Hydriodic 

Iodide 

HI 

Chloric 

Chlorate 

HC103 

(HO)C1V02 

C1205 

locdc 

lodate 

HI03 

(HO)IVO2 

I2O6 

THE  OPERATIONS  OF  ANALYSIS. 

§18.  Chemical  analysis  is  the  determination  of  any  or  all  of  the  compo- 
nents of  a  given  portion  of  matter,  whether  this  be  solid,  liquid  or  gaseous. 
A  portion  of  matter  is  made  up  of  one  or  more  definite  and  distinct  sub- 
stances, or  chemical  individuals,  each  of  which  is  either  a  "  compound  "  or 
an  "element  "  and  is  always  and  everywhere  the  same.  It  is  required  in 
analysis  to  detect  a  chemical  compound  as  a  body  distinct  from  the 
chemical  elements  that  have  formed  it.  For  example,  the  analyst  may 
have  in  hand  a  mixture  containing  sodium  sulphate,  Na  SO,  ;  sodium  sul- 
phite, Na,SO  i ,  and  sodium  thiosulphate,  Na2S.03 ,  but  not  containing  any 


14  THE  OPERATIONS  OF  ANALYSIS.  §19. 

sodium  or  sulphur  or  oxygen  as  these  bodies  are  severally  known  to  the 
world  and  described  in  chemistry.  In  this  instance  the  analyst  in  his 
ordinary  work  does  not  separate  the  sulphur  or  the  sodium,  as  elements 
uncombined  with  oxygen,  either  in  qualitative  or  in  quantitative  oper- 
ations. Each  one  of  the  compounds  of  the  sulphur  with  the  oxygen  is 
usually  sought  for  and  found  and  weighed  as  a  chemical  individual.  Cer- 
tain of  the  chemical  elements,  however,  are  frequently  separated  free  from 
all  combination,  as  a  method  of  determination  of  their  compounds. 

§19.  The  analysis  of  gaseous  material  is  termed  Gas  Analysis;  that  of 
mixtures  of  the  complex  compounds  of  carbon,  Organic  Analysis.  An 
examination  of  organic  matter,  when  limited  to  a  determination  of  its  ulti- 
mate chemical  elements  is  styled  Ultimate  Organic  Analysis.  When  it  is 
undertaken  to  determine  individual  carbon  compounds  actually  existing  in 
organic  matter,  it  has  been  spoken  of  as  Proximate  Organic  Analysis.  If 
the  same  distinction  were  to  be  applied  to  inorganic  analysis,  we  should 
have  to  say  that  it  is  mostly  "proximate"  but  is  sometimes  "ultimate" 
in  its  methods  of  operation. 

§20.  The  term  Qualitative  Chemical  Analysis  as  commonly  used  is  con- 
fined to  a  chemical  examination  of  material,  chiefly  inorganic,  in  the  solid 
or  liquid  state,  the  inquiry  being  limited  for  the  most  part  to  well  known 
substances. 

§21.  In  the  methods  of  analysis  of  a  mixture,  it  is  often  required  to 
separate  individual  substances  from  each  other,  but  sometimes  a  distinct 
compound  can  be  identified  and  sometimes  its  quantity  can  be  estimated 
while  it  is  in  the  presence  of  other  bodies.  Both  the  identification  and 
separation  are  accomplished,  nearly  always,  by  effecting  changes,  physical 
and  chemical. 

Methods  of  analysis  are  as  numerous  as  are  the  ways  of  bringing  into 
action  the  physical  and  chemical  forces  by  which  chemical  changes  are 
wrought.  The  characteristics  of  any  chemical  individual,  by  which  it  is 
distinguished  and  removed  from  others,  lie  in  its  responses  to  the  physical 
and  chemical  forces,  including  especially  the  chemical  action  of  certain 
well  known  compounds  called  reagents. 

§22.  The  response  toward  heat  and  pressure  fixes  the  melting  and  boiling 
points,  its  ordinary  solid  or  liquid  or  gaseous  state.  The  operations  "in 
the  dry  way  "  are  done  over  a  flame  or  in  a  furnace,  with  or  without  solid 
"reagents"  and  with  regard  to  oxidation.  They  represent  some  of  the 
methods  of  metallurgical  manufacture.  The  liquid  state,  whether  by 
fusing  or  by  solution,  is  the  state  commonly  necessary  or  favorable  to  chem- 
ical change  and  its  control. 

§23.  The  deportment  of  a  solid  substance  toward  light  comprises  its 
color  and  that  of  its  solutions,  as  well  as  that  of  its  vapor,  in  ordinary  light, 


§27.  THE   OPEHATIOXS   OF  ANALYMX.  15 

and  the  bands  and  primary  colors  it  exhibits  in  the  uses  of  the  spectroscope 
(Crookes,  J.  C.,  1889,  55,  255;  Welsbach,  M.,  1885,  6,  47). 

§24.  The  conduct  of  a  chemical  compound  in  electrolysis  is,  in  various 
cases,  a  means  both  of  identification  and  of  separation.  Electric  conduc- 
lirity  methods  are  used  for  establishing  the  presence  or  absence  of  minute 
traces  of  substances  (Kohlrausch  Whitney,  Z.  pJiys.  (77?..,  1896,  20,  44). 
Again,  traces  of  dissolved  matters  too  minute  for  other  means  of  detection 
can  be  revealed  by  the  difference  of  electric  potential  between  electrode  and 
solution  (Ostwald,  Lelirb.,  2  Aufl.,  II,  1,  881;  Behrend,  Z.  phys.  Cli.,  1893, 
11,  466;  Hulett,  Z.  phys.  Cli.,  1900,  33,  611). 

§25.  By  far  the  most  extensive  of  the  resources  of  analysis  lie  in  the 
chemical  reaction  of  one  definite  and  distinct  substance  with  another,  ac- 
cording to  the  character  of  each,  giving  rise  to  a  chemical  product  having 
peculiarities  of  its  own  in  evidence  of  its  origin.  In  this  way  the  com- 
pounds are  bound  in  regular  relations  to  each  other.  Therefore  it  belongs 
to  the  analyst  to  gain  personal  acquaintance  with  the  behavior  of  the  repre- 
sentative constituent  bases  and  acids  toward  each  other. 

§26.  Operations  for  chemical  change  are  commonly  conducted  in  solu- 
tion. The  material  for  analysis  is  dissolved,  and  is  treated  with  reagents 
that  are  in  solution.  A  solid  or  a  gas  is  dissolved  in  a  liquid  in  making  a 
solution.  When  the  dissolved  substance  is  converted  into  one  that  will 
not  dissolve  a  precipitate  is  formed.  It  is  necessary  therefore  to  under- 
stand the  nature  of  solution  and  to  give  heed  to  its  obvious  limitations. 
Certain  facts  and  conclusions  as  to  the  chemical  state  of  dissolved  com- 
pounds are  presented  under  the  head  next  following,  "  Solution  and  loniza- 
fcion."  But  it  must  first  be  observed  that  the  universal  solvent,  water,  is 
always  understood  to  be  present  in  somewhat  indefinite  proportion  in  opera- 
tions "  in  the  wet  way."  It  serves  as  a  vehicle,  as  such  not  being  included 
in  any  statement  of  the  substances  operated  upon,  nor  formulated  in  equa- 
tions, any  more  than  is  the  material  of  the  test  tube,  but  often  some  portion 
of  it  enters  into  combination  or  suffers  decomposition,  and  then  it  must  be 
placed  among  the  substances  engaged  in  chemical  change. 

§27.  No  other  property  of  substances  has  so  great  importance  in  analysis 
and  in  all  chemical  operations,  as  their  solubility  in  water.  It  must  never 
be  forgotten  that  there  are  degrees  of  solubility,  but  there  is  hardly  such  a 
fact  as  absolute  solubility,  or  insolubility,  regardless  of  the  proportion 
of  the  solvent.  There  are  liquids  which  are  miscible  with  each  other 
in  all  proportions,  but  solids  seldom  dissolve  in  all  proportions  of  the  sol- 
vent, neither  do  gases.  For  every  solid  or  gas,  there  is  a  least  quantity  of 
solvent  which  can  dissolve  it.  One  part  of  potassium  hydroxide  is  soluble 
in  one-half  part  of  water  (or  in  any  greater  quantity),  but  not  in  a  less 
quantity  of  the  solvent.  One  part  of  sodium  chloride  requires  at  least  two 


16  THE  OPERATIONS  OF  ANALYSIS.  §28. 

and  a  half  parts  of  water  to  dissolve  it.  One  part  of  mercuric  chloride  will 
dissolve  in  two  parts  of  water  at  100  degrees,  but  when  cooled  to  15  degrees 
so  much  of  the  salt  recrystallizes  from  the  solution,  that  it  needs  twelve 
parts  more  of  water  at  the  latter  temperature  to  keep  a  perfect  solution. 
Lead  chloride  dissolves  in  about  twenty  parts  of  hot  water,  about  half  of 
the  salt  separating  from  the  solution  when  cold.  Calcium  sulphate  dis- 
solves in  about  500  times  its  weight  of  water — this  dilute  solution  forming 
one  of  the  ordinary  reagents.  Barium  sulphate  is  one  of  the  least  soluble 
precipitates  obtained,  requiring  about  430,000  parts  of  water  for  its  solution 
at  ordinary  temperature  (Hollemann,  Z.  pliys.  Ch.,  1893,  12,  131).  In  ordi- 
nary reactions  it  is  not  appreciably  soluble  in  water.  Lead  sulphate  dis- 
solves in  about  21,000  parts  of  water:  in  many  operations  this  solubility 
may  be  disregarded,  but  in  quantitative  analysis  the  precipitate  is  washed 
with  alcohol  instead  of  water,  losing  less  weight  with  the  former  solvent. 
These  examples  indicate  the  necessity  of  discriminating  between  degrees  of 
solubility.  Also  the  solubility  of  a  particular  compound  is  dependent  upon 
the  physical  form  of  that  compound  (§69,  5  &);  e.  g.,  amorphous  magnesium 
ammonium  phosphate  is  quite  soluble  in  water,  the  crystalline  salt  being 
almost  insoluble.  The  solubility  of  a  solid  is  ,-ilso  dependent  upon  the  size 
of  the  particles  of  the  solid,  a  finely  divided  solid  being  more  soluble  than 
large  particles  of  the  same  substance  (Hullett  and  Allen,  J.  Am.  Soc.,  24, 
667,  1902).  In  analysis  it  is  customary  to  heat  and  then  allow  a  precipitate 
to  stand  with  the  solution  in  which  it  was  formed  in  order  to  obtain  com- 
plete precipitation.  When  a  solvent  has  dissolved  all  of  a  substance  that  it 
can  at  a  particular  temperature,  in  contact  with  the  solid,  the  solution  is 
said  to  be  saturated  at  that  temperature.  It  frequently  happens  that  a 
saturated  solution  of  a  substance  at  a  higher  temperature  may  be  cooled 
without  separation  of  the  solid.  Such  a  solution  (at  the  lower  temperature) 
is  said  to  be  supersaturated  and  precipitation  frequently  is  induced  by 
jarring  the  solution,  more  surely  by  adding  a  crystal  of  the  dissolved 
substance. 

§28.  The  ordinary  liquid  reagents  are  water  solutions — concentrated  sul- 
phuric acid  and  carbon  disulphide  being  exceptions.  Hydrochloric  acid, 
liquid  hydrogen  sulphide,  and  ammonium  hydroxide  are  solutions  of  gases 
in  water;  on  exposure  to  the  air  these  gases  gradually  separate  from  their 
solutions.  All  these  gases  escape  much  more  rapidly  when  their  solutions 
are  warmed.  The  majority  of  liquid  reagents  are  solids  in  aqueous  solu- 
tion. (See  the  list  of  Reagents.) 

§29.  Substances  are  said  to  dissolve  in  acids,  or  in  alkalis,  and  this  is 
termed  chemical  solution;  more  definitively  it  is  chemical  action  and  solu- 
tion, the  solution  being  counted  as  a  physical  change.  We  say  that  cal- 
cium oxide  dissolves  (chemically)  in  hydrochloric  acid;  that  is,  in  the 


§33.  THE  OPERATIONS  OF  ANALYSIS.  17 

reagent  named  hydrochloric  acid,  a  mixture  of  that  acid  and  water.  The 
acid  unites  with  the  calcium  oxide,  forming  a  soluble  solid,  which  the  water 
dissolves.  Absolute  hydrochloric  acid  cannot  dissolve  calcium  oxide. 

§30.  Solids  can  be  obtained,  without  chemical  change,  from  their  aqueous 
solutions:  Firstly,  by  evaporation  of  the  water.  This  is  done  by  a  careful 
application  of  heat.  Secondly,  solids  can  be  removed  from  solution,  with- 
out chemical  change,  by  (physical)  precipitation — accomplished  by  modify- 
ing the  solvent.  If  a  solution  of  potassium  carbonate,  or  of  ferrous  sul- 
phate, be  dropped  into  alcohol,  a  precipitate  is  obtained,  because  the  salts 
will  not  dissolve,  or  remain  dissolved,  in  the  mixture  of  alcohol  and  water. 
But,  in  analysis,  precipitation  is  more  often  effected  by  changing  the  dis- 
solved substance  instead  of  the  solvent. 

§31.  Solids  can  be  separated  from  their  solution  by  precipitation  due  to 
chemical  change,  to  the  extent  that  the  product  is  insoluble  in  the  quantity 
of  the  solvent  present.  Calcium  can  be  in  part  precipitated  from  not  too 
dilute  solutions  of  its  salts,  by  addition  of  sulphuric  acid;  but  there  still 
remains  not  precipitated  the  amount  of  calcium  sulphate  soluble  in  the 
water  and  acid  present,  which  is  enough  to  give  an  abundant  precipitate 
with  ammonium  oxalate,  the  precipitated  sulphate  being  previously  re- 
moved by  filtration. 

Time  and  heat  are  required  for  the  completion  of  most  precipita- 
tions. If  it  is  necessary  to  remove  a  substance,  by  precipitation,  before 
testing  for  another  substance,  the  mixture  should  be  warmed  and  allowed 
to  stand  for  some  time,  before  filtration.  Neglect  of  these  precautions  often 
occasions  a  double  failure;  the  true  indication  is  lost,  and  a  false  indication 
is  obtained. 

§32.  Eeagents  should  be  added  in  very  small  portions,  generally  drop  by 
drop.  Often  the  first  drop  is  enough.  Sometimes  the  precipitate  redis- 
solves  in  the  reagent  that  produced  it,  and  this  is  ascertained  if  the  reagent 
be  added  in  small  portions,  with  observation  of  the  result  of  each  addition. 
If  it  is  a  final  test,  a  quantity  of  precipitate  which  is  clearly  visible  is  suffi- 
cient, but  if  the  precipitate  is  to  be  filtered  out  and  dissolved,  a  considerable 
quantity  should  be  formed.  If  the  precipitate  is  to  be  removed  and  the 
filtrate  tested  further,  the  precipitation  must  be  completed — by  adding  the 
reagent  as  long  as  the  precipitate  increases,  with  the  warmth  and  time 
requisite  in  the  operation;  and  a  drop  of  the  same  reagent  should  be  added 
to  the  filtrate  to  obtain  assurance  that  the  precipitation  has  been  completed. 
It  will  be  found,  with  a  little  experience,  that  some  reagents  must  be  used 
in  relatively  large  quantities.  On  the  contrary,  the  acids,  sulphuric,  hydro- 
chloric and  nitric,  are  required  in  a  volume  relatively  very  small. 

§33.  Certain  very  exact  methods  of  identification  can  be  conducted  by 
drop  tests  upon  a  black  or  white  ground,  or  upon  a  glass  slide  and  especially 


THE   OPERATION*   OF  ANALYSIS.  §34. 

with  the  help  of  a  microscope  and  with  studies  of  crystalline  form.  Further 
see  Behrens,  Z.  1891,  30,  1^5;  and  Herrnschmidt  and  Capelle,  Z.  189;  .  32, 
608. 

§34.  Precipitates  are  removed — usually  hy  filtration,  sometimes  by  d  ^can- 
tation.  If  they  are  to  be  dissolved,  they  must  he  first  washed  till  free  from 
all  the  substances  in  solution.  For  complete  precipitation  some  excess  of 
the  reagent  must  have  "been  used.  Beside  the  reagent  there  are  othe  dis- 
solved matters,  after  precipitations,  some  of  which  are  indicated  Ir  the 
equation  written  for  the  change.  All  these  dissolved  substances  pemeate 
and  adhere  to  the  porous  precipitate  with  greater  or  less  tenacity.  If  they 
are  not  wholly  washed  away,  some  portion  of  them  will  he  mixed  with  the 
dissolved  precipitate.  Then,  the  separation  of  substances,  the  only  object 
of  the  precipitation,  is  not  accomplished,  while  the  operator,  proceeding 
just  as  though  it  was  accomplished,  undertakes  to  identify  the  members  of 
a  group  by  reactions  on  a  mixture  of  groups.  The  washing,  on  the  :ilter. 
is  best  completed  by  repeated  additions  of  small  portions  of  water — around 
the  filter  border,  from  the  wash  bottle — allowing  each  portion -to  pass 
through  before  another  is  added.  The  washings  should  be  tested,  from 
time  to  time,  until  they  are  free  from  dissolved  substances. 

§35.  In  dissolving  precipitates — by  aid.  of  acids  or  other  agents — use 
the  least  possible  excess  of  the  solvent.  Fndeavor  to  obtain  a  solation 
nearly  or  quite  saturated,  chemically.  If  a  large  excess  of  acid  is  carried 
into  the  solution  to  be  operated  upon,  it  usually  has  to  be  neutralized,  and 
the  solution  then  becomes  so  greatly  encumbered  and  diluted  that  reactions 
become  faint  or  inappreciable.  Precipitates  may  be  dissolved  on  the  filter, 
without  excess  of  solvent,  by  passing  iho  panic  portion  of  the  (diluted) 
solvent  repeatedly  through  the  filter,  following  it  once  or  twice  with  a  few 
drops  of  water.  The  mineral  acids  should  be  diluted  to  the  extent  required 
in  each  case.  For  solution  of  small  quantities  of  carbonates  and  ?omo 
other  easily  soluble  precipitates  the  acids  may  be  diluted  with  fifty  times 
their  weight  of  water.  Washed  precipitates  may  also  be  dissolved  ri  the 
test-tube,  by  rinsing  them  from  the  filter,  through  a  puncture  made  in  its 
point,  with  a  very  little  water.  If  the  filter  be  wetted  before  filtration,  the 
precipitate  will  not  adhere  to  it  so  closely. 

§36.  When  a  reagent  is  added  in  order  to  produce  a  change  in  the  acid, 
alkaline  or  neutral  condition  of  the  solution,  the  addition  of  suff.cient 
reagent  to  cause  the  desired  change  should  always  be  governed  by  testing 
a  drop  of  the  solution,  on  a  glass  rod.  with  a  piece  of  litmus  paper. 

£37.  AY  hen  substances  in  separate  solution  are  brought  together,  an 
evidence  of  the  formation  of  a  new  substance  is  the  appearance  of  a  solid 
in  the  mixture,  i.e.  aprecipitate.  A  chemical  change  between  dissolvec  sub- 
stances— salts,  acids,  and  bases — will  be  practically  complete  when  one  or 


§40.  THE  OPERATIONS  OF  ANALYSIS.  19 

more  of  the  products  of  such  change  is  a  solid  or  a  gas,  not  soluble  in  the 
mixtiTC.  As  an  example,  Calcium  carbonate  +  Hydrochloric  acid  —  Cal- 
*•  in  in  chloride  +  Water  -(-  Carbon- dioxide  (gas). 

§3'.  Tn  the  practice  of  qualitative  analysis,  the  student  necessarily  refers 
to  authority  for  the  composition  of  precipitates  and  other  products.  For 
exam  tie,  when  the  solution  of  a  carbonate  is  added  to  the  solution  of  a 
calcivm  sail,  a  precipitate  is  obtained;  and  it  has  been  ascertained  by  quanti- 
tativ  analysis  that  this  precipitate  is  normal  calcium  carbonate,  CaCO, , 
invar 'ably.  Were  there  no  authorized  statement  of  the  composition  of  this 
precipitate,  the  student  would  be  unable,  without  making  a  quantitative 
anah  ?is,  to  declare  its  formula,  or  to  write  the  equation  for  its  production. 
Whe:.  the  results  of  analytical  operations  are  substances  of  unknown,  uncer- 
tain, or  variable  composition,  equations  cannot  be  given  for  them. 

g3l\  The  written  equation  represents  only  the  substances,  and  the  quan- 
tity (>f  each,  which  actually  undergo  the  chemical  change  that  is  to  be 
expressed.  Thus,  if  a  reagent  is  used  to  effect  complete  precipitation,  an 
exces -;  of  it  must  be  employed,  beyond  the  ratio  of  its  combining  weight  in 
the  e 'illation.  That  is,  if  magnesium  sulphate  be  employed  to  precipitate 
bariu  ii  chloride,  the  exact  relative  amount  of  magnesium  sulphate  indicated 
by  th  3  equation:  BaCL  +  MgS04  =  BaS04  -f-  MgCl, ,  fails  to  precipitate  all 
of  th  ?  barium.  The  soluble  sulphate  must  be  in  a^slight  excess.  On  the 
other  hand,  to  effect  complete  precipitation  of  the  sulphate  the  barium 
must  be  in  a  slight  excess. 

§40.  By  translating  chemical  equations  into  statements  of  proportional 
parts  by  weight,  they  are  prepared  to  serve  as  standard  data  of  absolutely 
pure  materials,  and  applicable  in  operations  of  manufacture,  with  large  or 
small  quantities,  after  making  due  allowance  for  moisture  and  other  im- 
purities, necessary  excess,  etc.  In  quantitative  analysis  the  equation  is  the 
constant  reliance.  For  example,  in  dissolving  iron  by  the  aid  of  hydro- 
chloric acid,  we  have  the  equation: 

Fe  +  2HC1  =  FeCL  +  H2  . 

55.8  +  72.9  =  126.8  +  2  . 

Also  in  precipitating  ferrous  chloride  by  sodium  phosphate,  we  have  the 
equation: 

FeCL  +  BTa2HP04,12H20  =  FeHPO4  +  2NaCl  +  12H20  . 

126.8  +  (142.2  +  216)  =  151.8  +  117  . 

Suppose  it  is  desired  to  determine  from  the  above: 

(1)  How  much  hydrochloric  acid,  strength  32  per  cent,  is  required  to 
dissolve  100  parts  of  iron  wire. 

(2)  What  quantities  of  32  per  cent  hydrochloric  acid  and  iron  wire  are 
necessary  to  use  in  preparing  100  parts  of  absolute  ferrous  chloride. 


20  SOLUTION  AND  IONIZATION.  §41, 

(3)  What  materials  and  what  quantities  of  them,  may  be  used  in  prepar- 
ing 100  parts  of  ferrous  phosphate. 

In  practice  allowance  must  be  made  for  the  facts  that  the  iron  wire  will 
not  be  quite  pure,  and  that  a  considerable  excess  of  the  hydrochloric  acid 
would  be  necessary  for  the  complete  solution  of  the  iron.  Also  that  some 
excess  of  the  phosphate  would  be  necessary  to  the  full  precipitation  of  the 
iron.  Irrespective  of  impurities,  oxidation  product  and  excess,  the  re- 
quired quantities  are  found  by  the  combining  weights  as  follows: 

j  55.8/72.9  =  100/x  =  parts  of  absolute  HC1  for  100  parts  of  iron  wire. 
'  (  32/100  =x/y  =  parts  of  32  per  cent  HC1  for  100  parts  of  iron  wire. 

f  126.8/72.9=  100/x. 

2.  -j  32/100  =x/y  =  parts  of  32  per  cent  HC1  for  100  parts  of  FeCl2 ,  absolute, 
(  129.8/55.8  =  100/z  =  parts  of  iron  wire  for  100  parts  of  FeCl2 . 
,  151.8/72.9  =  100/x. 

„  j  32/100  =x/y  =  parts  of  32  per  cent  HC1  for  100  parts  of  FeHPO4 . 
1  151.8/55  8  =100/z  =  parts  of  metallic  iron  for  100  parts  of  FeHPO4 . 
'  151.8/358.2  =  100/u  =  parts  of  Na>HPO4,  12H:O  for  100  parts  of  FeHPO4. 

Practice  in  reducing  the  combining  numbers  of  the  terms  in  an  equation 
to  simple  parts  by  weight,  is  a  very  instructive  exercise,  even  in  the  early 
part  of  qualitative  chemistry.  It  enforces  correct  and  clear  ideas  of  the 
significance  of  formula  and  equations,  and  refers  all  chemical  expressions 
to  the  facts  of  quantitative  work. 

§41.  The  chief  requirement  in  qualitative  practice  is  an  experimental 
acquaintance  with  the  chemical  relations  of  substances,  rather  than  the 
identification  of  one  after  the  other  by  routine  methods.  The  acids  and 
bases,  the  oxidizing  and  reducing  agents,  are  all  linked  together  in  a  net- 
work of  relations,  and  the  ability  to  identify  one,  as  it  may  be  presented  in 
any  combination  or  mixture,  depends  upon  acquaintance  with  the  entire 
fraternity. 

§42.  The  full  text  of  the  book,  rather  than  the  analytical  tables,  should  be 
taken  as  the  guide  in  qualitative  operations,  especially  in  those  upon  known 
material.  The  tabular  comparisons  are  commended  to  attention,  especially 
for  review.  In  actual  analysis,  the  tables  serve  mainly  as  an  index  to  the 
body  of  the  work. 

SOLUTION  AND  IOKIZATION. 

§43.  The  Theory  of  Electrolytic  Dissociation,  proposed  by  Arrhenius  in 
1887  (Z.  pliys.  Ch.,  1887,  1,  631),  assumes  that  acids,  bases  and  salts  in 
water  solution  are  present  not  as  the  intact  molecule  but  split  up  into  two 
or  more  parts  called  ions  which  are  charged  with  negative  or  positive  elec- 
tricity. The  facts  upon  which  the  theory  is  based  are  the  osmotic  pressure,* 

*  The  pressure  by  virtue  of  which  a  soluble  substance  in  contact  with  the  solvent,  as  common 
salt  in  water,  is  enabled  to  rise  against  the  force  of  gravity  and  distribute  itself  uniformly  through- 
out the  solvent,  just  as  gas  by  the  virtue  of  the  gas-pressure  occupies  the  entire  space  at  its  disposal. 


§43.  SOLUTION  AND  IONIZAT1ON.  21 

lowering  of  the  freezing  point,  raising  of  the  boiling  point  and  electrical 
conductivity  of  water  solutions  of  acids,  bases  and  sails.  Because  such 
water  solutions  conduct  the  electric  current,  acids,  bases  and  salts  are 
culled  electrolytes. 

The  osmotic  pressure  of  a  solution  is  believed  to  be  proportional  to  the 
number  of  particles  of  the  dissolved  substance  present  in  unit  volume  of 
the  solution.  In  the  case  of  non-electrolytes  the  osmotic  pressure  is  pro- 
portional to  the  molecular  weight;  but  the  osmotic  pressure  of  electrolytes 
is  greater  than  corresponds  to  their  molecular  weight.  This  is  readily  ex- 
plained by  the  assumption  that  some  of  the  molecules  are  split  into  two  or 
more  parts. 

In  a  similar  manner  the  freezing  point  of  a  liquid  is  lowered  by  the  pres- 
ence of  a  dissolved  substance,  and  the  amount  of  the  lowering  is  propor- 
tional to  the  number  of  dissolved  particles.  In  the  case  of  non-electrolytes 
the  lowering  is  proportional  to  the  molecular  weight,  but  dissolved  electro- 
lytes depress  the  freezing  point  to  a  greater  extent,  indicating  dissociation 
of  the  molecules.  The  boiling  point  of  a  liquid  is  raised  by  the  presence  of 
a  dissolved  substance,  but  this  effect  is  greater  in  the  case  of  electrolytes 
because  the  molecule  is  dissociated.  This  reasoning  is  similar  to  that 
applied  to  gas  pressures. 

The  gas-laws  (Boyle's,  Guy-Lussac'a,  Henry's,  and  Dalton's)  are  found 
to  hold  for  dissolved  substances,  osmotic  pressure  being  substituted  for  gas- 
pressure  (van  't  Hoff,  Z.  phys.  Ch.,  1887,  1,  481;  Morse  and  Frazer,  Am.9 
34,  1  (1905);  37,  324,  425,  558;  38,  175  (1907);  Lewis,  J.  Am.  Soc.,ZQ, 
6G8  (1908) ).  Avogadro's  Hypothesis  is  therefore  applicable  to  solutions  as 
well  as  to  gases,  and  as  abnormal  gas-pressure  points  to  dissociation  in  the  gas 
(NH4C1 ,  PC16)  so  excessive  osmotic  pressure,  lowering  of  freezing  point  and 
raising  of  boiling  point  is  taken  as  indicating  dissociation  of  the  dissolved 
substance.  The  osmotic  pressure  as  well  as  the  abnormal  freezing  and 
boiling  point  may  be  taken  as  a  measure  of  this  dissociation. 

The  fact  that  solutions  of  non-  electrolytes  do  not  conduct  the  electric 
current  while  solutions  of  electrolytes  do  conduct  the  electric  current  indi- 
cates that  molecules  are  incapable  of  carrying  the  current,  but  that  the 
component  parts  into  which  the  molecule  is  split  carry  the  current. 

Faraday  gave  the  name  ions  to  the  components  of  a  substance  conducting 
the  electric  current  in  solution.  It  is  an  observed  fact  that  transmission 
of  the  current  by  a  solution  is  always  accompanied  by  movement  of  the 
ions  in  opposite  directions  (Hittorf,  Pogg.  1853,  89,  177).  This  is  quite 
independent  of  any  separations  taking  place  at  the  electrodes.  From  this 
it  is  concluded  that  the  ions  carry  the  electricity  from  one  pole  to  the 
other  through  the  solution.  If  the  ions  are  the  carriers  of  electricity  then 
the  power  of  a  solution  to  conduct  the  current  will  be  in  proportion  to  their 


£2  SOLUTION  AND  JONIZAT10N.  §43. 

number,  that  is,  to  the  extent  of  dissociation  of  the  dissolved  substance. 
And  experiment  shows  that  the  dissociation  calculated  from  the  osmotic 
pressure  is  identical  with  the  dissociation  calculated  from  the  electric 
conductivity. 

Further,  if  in  analysis  of  a  substance  in  solution  we  are  dealing  not  with 
the  substance  in  its  integrity  but  with  certain  ions,  then  our  ordinary 
analytical  reactions  are  reactions  of  the  ions,  and  we  may  expect  that  where 
the  substance  for  some  reason  is  transformed  from  the  ionized  condition 
to  the  undivided  molecule  these  reactions  will  fail.  Here  again  the  chemi- 
cal activity  will  be  proportional  to  the  number  of  ions;  and  experiment 
shows  that  quantitative  parallelism  exists,  to  take  the  case  of  acids,  between 
(1)  the  characteristic  acid  activity — the  dissolving  of  metals,  the  influence 
as  catalyzer  on  such  changes  as  the  inversion  of  cane-sugar  and  the  saponi- 
fication  of  esters;  (2)  the  extent  of  dissociation  as  indicated  by  osmotic 
pressure,  and  (3)  the  extent  of  dissociation  as  indicated  by  electric  conduct- 
ivity. The  same  parallelism  holds  for  other  bodies  in  solution.  A  water 
solution  of  an  acid  such  as  hydrochloric  or  sulphuric  should  not  be  regirded 
as  consisting  of  molecules  of  the  acid  and  water  but  as  a  solution  of  mole- 
cules of  HC1  and  the  ions  H  and  Cl  or  molecules  of  H,S04  and  ions  H  ai:il  S04 
respectively,  the  hydrogen  ion  being  charged  with  positive  electricity  ai  d  the 
acid  ion  with  negative  electricity.  The  very  active  acids  and  bases  ai.d  the 
neutral  salts  undergo  wide  dissociation  in  water  solution,  while  weak  acids 
and  bases  retain  almost  entirely  the  non-dissociated  condition,  the  strength 
of  the  acids  being  proportional  to  the  concentration  of  the  H  ion. 

The  Electrolytic  Dissociation  Theory  in  its  assumption  of  a  separation 
into  ions  groups  together  and  gives  system  and  meaning  to  these  three 
classes  of  facts,  experimentally  absolutely  independent  and  up  to  Arrhonius' 
time  without  any  suspected  relationship.  In  each  case  the  results  calculated 
on  the  assumption  of  such  a  dissociation  are  in  quantitative  agreement  with 
those  obtained  by  measurement. 

Corresponding  in  actual  experience  to  the  view  that  the  common  analyti- 
cal reactions  are  due  to  the  ions  rather  than  to  the  molecule  as  a  whole,  is 
the  analyst's  practice  of  testing  for  acid  radicle  or  basic  radicle  without 
regard  to  the  other  component.  For  instance,  HJS  or  K2S  will  produce 
precipitates  of  metallic  sulphides  because  the  sulphur  is  present  in  solu- 
tion as  an  ion.  On  the  other  hand  HSO;,  H2  S03 ,  or  H,S_03  wi  1  not 
precipitate  metals  as  sulphides,  because  in  these  acids  the  ion  is  com- 
posed of  the  sulphur  and  the  oxygen  present.  Further,  HgCL  in  its 
chemical  behavior  is  unlike  other  mercuric  salts  and  unlike  other  chlorides. 
The  mercury  is  not  readily  precipitated  by  alkali  hydroxides  n-n*  is 
the  chloride  readily  precipitated  by  silver  salts.  In  agreement  with  this, 
its  conductivity  and  osmotic  pressure  are  also  unlike  those  of  the  great 


§44.  SOLUTION  AND  lONtZATIONn  23 

majority  of  neutral  salts,  both  pointing  to  very  slight  dissociation  into  the 
ions.  CdClo  is  another  neutral  salt  anomalous  in  that  its  conductivity  and 
osmc:ic  pressure  are  both  low.  And  here  also  for  precipitation  of  the 
chlor'.de  a  considerable  concentration  of  the  reagent  is  necessary.  Similar 
instances  of  the  parallelism  referred  to  are  numberless. 

§40:.  The  Law  of  Mass-Action  embodies  the  familiar  principle  that  the 
chemical  activity  of  a  substance  is  proportional  to  its  concentration.  It 
was  irst  recognized,  although  imperfectly,  by  Berthollet  and  was  given 
math  3inatical  expression  by  Guldberg  and  Waagc  in  1867.  The  latter 
investigators  found  it  to  accord  well  with  the  observed  facts  in  some  cases; 
in  others  there  were  wide  discrepancies  which  were  later  shown  by  Ar- 
rhenuis  to  disappear  when  the  concentration,  not  of  the  reacting  body  as  a 
whole  but  only  of  that  part  present  in  the  ionized  condition,  was  taken 
into  consideration.  We  must  assume  that  every  chemical  reaction  is  rever- 
sible, that  is,  that  none  of  them  proceed  until  the  reacting  substances  are 
completely  transformed.  Then  by  a  simple  process  of  reasoning  it  is  found 
that  \rhen  equilibrium  sets  in  the  product  obtained  by  multiplying  together 
the  concentrations  of  the  reacting  substances  will  be  in  a  certain  definite 
ratio  to  the  product  of  the  concentrations  of  the  substances  formed,  con- 
centration being  defined  as  the  quantity  in  unit  volume.*  For  example, 
in  fie  reaction  indicated  by  the  equation  CH,CO,H  -\-  C2H5OH  — 
CH3C  0,0,11,  -f  H20  ,  when  equilibrium  sets  in  ab  =  kcd ,  in  which  a  and  b 
are  t-ie  concentrations  of  acid  and  alcohol  respectively,  c  and  d  those  of 
ester  and  water,  while  k  is  a  constant  peculiar  to  the  reaction.  Where  the 
reaction  is  a  dissociation,  as  with  gaseous  NH4C1 ,  we  have  ab  =  k'c  ,  a  and  b 
repre  renting  the  concentrations  of  NH3  and  HC1  respectively,  c  that  of  the 
uncle  jomposed  NH4C1 ,  and  k'  the  constant  characteristic  of  this  change. 
Dissociation  into  ions  must  follow  the  same  laws,  and  for  the  electrolytic 
dissociation  of  acetic  acid  a  similar  equation  holds,  a  and  b  in  this  case 
standing  for  concentration  of  H  and  acetic  ions,  c  for  concentration  of  non- 
dissociated  acetic  acid,  while  the  constant  is  one  governing  only  this  par- 
ticular dissociation.  It  is  apparent  from  each  of  these  equations  that,  if 
we  aod  one  of  the  products  of  the  reaction  and  thus  increase  its  concentra- 
tion, the  concentration  of  the  other  product  must  decrease  in  the  same 
proportion — the  extent  of  the  reaction  will  be  decreased;  while,  on  the 
other  hand,  removing  either  or  both  of  the  products  will  tend  to  make  the 
transformation  complete.  This  deduction  is  of  great  significance.  In 
making  ethyl  acetate  from  the  acid  and  alcohol,  in  order  to  use  the  materials 
as  completely  as  possible,  the  ester  is  distilled  off  as  rapidly  as  produced 

*  Tl  3  unit  of  quantity  is  the  molecular  weight  taken  in  grams  (the  "  mol  ")•  Where  there  are 
18.23  |  rams  HCI  in  a  liter  either  in  solution  or  as  gas  the  concentration  is  M,  where  there. are 
72.92  ^;ams  in  the  same  volume  the  corcentratiori  is  2  and  so  on, 


24  SOLUTION  AND  IONIZATION.  §45. 

\vhile  the  water  is  taken  up  by  some  absorbent.  Introducing  gaseous  NH:j 
or  HC1  diminishes  the  dissociation  of  NH4C1  by  heat,  and  similarly  adding 
either  H  ions  or  acetic  ions  will  diminish  the  dissociation  of  acetic  acid. 
Aeetic  acid  is  much  weakened  by  the  presence  of  a  neutral  acetate.  A 
ferrous  solution  moderately  acidified  with  acetic  acid  gives  no  precipitate 
on  saturation  with  H2S,  but  on  addition  of  sodium  acetate  the  black  FeS 
is  brought  down.  Similarly  a  weak  base,  as  NH4OH  ,  is  made  still  less 
effective  by  the  presence  of  its  strongly-dissociated  neutral  salt,  as  NH4C1 . 
Quantitative  agreement  is  obtained  between  observed  effect  of  NH4C1  on 
NH4OH  as  saponifying  agent  and  that  calculated  from  the  equation: 

°NH  '  COH'  ~  k°NH  OH  (Arrhenius,  Z.  phys.  Cli.,  1887,  1,  110). 

In  general  every  acid  is  weakened  by  the  addition  of  the  neutral  salt  of 
the  acid  to  its  solution.  Similarly  bases  are  weakened  by  the  addition  of 
the  neutral  salt  of  the  base  to  its  solution. 

§45.  The  Solubility-Product. — In  the  saturated  solution  which  always 
remains  after  precipitation  we  have  the  usual  dissociation  equilibrium,  as: 

CAe  '  CC1'  ~~  AffCl '  ^ow  ^ne  quaT1tity  °^  non-dissociated  substance  in 
a  saturated  solution  is  invariable  and  the  right  side  of  this  equation  is 
therefore  constant.  That  is,  in  saturated  solution  the  product  of  the  con- 
centrations of  the  ions  is  always  the  same  for  a  given  substance  (Nernst). 
This  Ostwald  has  called  the  Solubility-Product.  Where  the  saturated  solu- 
tion is  made  by  bringing  the  salt  into  contact  with  the  solvent  c^ff  -  ~ "  CQ^  • 

From  such  a  solution  precipitation  will  take  place  on  addition  of  either  a 
silver  salt  or  a  chloride,  for  such  addition  largely  increases  the  concentration 
of  one  ion  and,  to  restore  equilibrium,  the  concentration  of  the  other  ion 
must  decrease  in  the  same  proportion,  which  is  possible  only  by  precipita- 
tion. From  this  follows  the  old  empirical  rule  to  add  an  excess  of  the 
reagent  in  making  a  precipitation.  Experiments  on  this  point  give  quanti- 
tative agreement  with  the  theory  (Nernst,  Z.  phys.  Ch.,  1889,  4,  372; 
Noyes,  Z.  phys.  Ch.,  1890,  6,  241;  1892,  9,  603;  1898,  26,  152). 

The      Solubility-Product      of      the      alkaline-earth      carbonates      is 

°M  "  CCO  "  ^n  ^e  soluti°n  of  a  neutral  salt,  as  CaCl2 ,  Ca  ions  are 

present  in  large  concentration.  When  a  substance  containing  C03  ions  in 
large  concentration  is  added,  as  Na2C03 ,  the  solubility-product  is  exceeded 
and  precipitation  takes  place.  Carbonic  acid,  however,  is  shown  by  con- 
ductivity and  osmotic  pressure  measurements  to  be  but  slightly  disso 
ciated,  that  is,  it  contains  few  C03  ions,  and  in  accord  with  this  is  the 
familiar  fact  that  the  alkaline  earths  are  not  precipitated  by  carbonic  acid. 
Similarly  the  fixed  alkali  hydroxides,  strongly  dissociated,  will  precipitate 


§46.  ORDER    OF  LABORATORY  STUDY.  25 

alkaline-earth  hydroxides,  while  ammonium  hydroxide,  shown  by  othei 
measurements  to  contain  but  few  hydroxyl  ions,  will  not. 

For   the   metallic   sulphides   the    solubility-product   is      j^-  ••    g" 

The  alkali  sulphides  as  normal  salts  contain  the  S  ion  in  large  concentra- 
tion and  so  produce  precipitation  even  of  the  more  soluble  sulphides  of 
the  Iron  and  Zinc  Groups.  The  slightly  dissociated  H2S  contains  sufficient 
S  ions  to  reach  the  solubility-product  of  the  sulphides  of  the  Silver,  Tin, 
and  Copper  Groups,  but  not  enough  to  attain  to  the  larger  solubility- 
product  of  the  Iron  and  Zinc  Group  sulphides.  A  strong  acid,  as  HC1 . 
containing  as  it  does  H  ions,  one  of  the  dissociation  products  of  H2S ,  drives 
back  the  dissociation  of  the  H2S,  so  decreasing  the  concentration  of  the 
S  ions  and  making  precipitation  of  the  sulphide  more  difficult. 

For  the  application  of  the  dissociation  theory  to  the  details  of  analytical 
work  we  are  indebted  chiefly  to  Ostwald.  See  his  "  Scientific  Foundations 
of  Analytical  Chemistry  "  and  "  Outlines  of  General  Chemistry/' 

ORDER  OF  LABORATORY  STUDY. 

846.  The  following  is  a  suggestive  outline  to  be  modified  by  the  teacher 
to  suit  the  ability  of  the  students,  and  the  amount  of  time  to  be  given  to 
the  study : 

a.  A  review  of  chemical  notation  and  the  writing  of  the  formulas  of  salts. 

&.  A  study  of  the  action  of  the  Fixed  Alkalis  upon  solutions  of  the  salts 
of  the  metals  in  the  order  of  their  groupings;  including  the  action  of  an 
excess  of  the  reagent.  The  fact  of  the  reaction  should  be  stated;  e.  </., 
the  addition  of  potassium  hydroxide  to  lead  acetate  produces  a  white  pre- 
cipitate readily  soluble  in  excess  of  the  reagent.  The  text  should  then  be 
consulted  for  the  products  of  the  reaction  (§57,  6«),  and  the  reactions  ex- 
pressed in  the  form  of  equations: 

2Pb(C2H302)2  -f-  4KOH  =  Pb2Q(OH),  *  (white)  +  4KC2H3O2  +  H2O 
FbjO(OH)2  -f  4KOH  (excess)  =  2K2PbO2  +  3H>O 
or  Pb(C2H3O2)il  +  4KOH  (excess)  =  K2PbO2  +  2KC2H3O2  +  2H2O . 

The  results  should  all  be  tabulated  and  then  summarized  in  the  form  of  a 
carefully  worded  generalization  (§205,  6a). 

c.  Action  of  Ammonium  Hydroxide  (volatile  alkali)  upon  solutions  of 
the  suits  of  the  metals,  etc.,  as  in  (b)  above;  e.g.,  the  addition  of  ammonium 
hydroxide  to  lead  nitrate  produces  a  white  precipitate  not  dissolving  in  ex* 
cess.  Consult  text  (§57,  Go)  and  write  the  equation: 

3Pb(N03)2  +  4NH4OH  =  2PbO  Pb(NO3)2  +  4NH4NO3  +  2H2O  . 
After   the   work  has   been   completed   in   the  laboratory  and  the  results 
*  It  Las  been  found  helpful  to  require  students  to  underscore  alt  precipitates. 


26  ORDER   OF  LABORATORY  STUDY.  §46. 

discussed  in  the  class  room.,  summarize  in  the  form  of  a  generalized  state- 
ment (§207,  6ff). 

d.  A  study  of  the  action  of  the  Fixed  Alkali  Carbonates,  and  generaliza- 
tion of  the  results  (§205,  6a). 

e.  A  study  of  the  action  of  Ammonium  Carbonate.     Summarize  the  re- 
sults (§207,  6a). 

f.  A  study  of  the  solvent  action  of  acids,  HC1 ,  HNOS ,  and  H2S04 ,  upon 
the  Hydroxides  and  Carbonates  obtained  by  precipitation. 

g.  Action  of  Hydrosulphuric  Acid  as  a  precipitating  agent  upon  salts  of 
the  metals  in  neutral  and  acid  solutions. 

h.  The  use  of  Ammonium  Sulphide  as  a  reagent. 

t.  The  solvent  action  of  acids,  HC1 ,  HN03,  and  HC2H,02 ,  upon  the 
sulphides  obtained  by  precipitation. 

;'.  Action  of  Hydrochloric  Acid  and  Soluble  Chlorides. 
Action  of  Hydrobromic  Acid  and  Soluble  Bromides. 
Action  of  Hydriodic  Acid  and  Soluble  Iodides. 

Jc.  Precipitation  by  Soluble  Sulphates,  Phosphates,  and  Oxalates. 

I.  The  solvent  action  of  Hydrochloric  and  Acetic  Acids  upon  the  Phos- 
phates obtained  by  precipitation. 

m.  The  reverse  of  certain  of  the  above  reactions  as  illustrating  the 
precipitation  of  Acids;  e.  g.,  the  addition  of  calcium  chloride  to  ammonium 
oxidate  produces  a  white  precipitate.  Consult  the  text  (§227,  8)  and  write 
the  equation:  (NH4)C204  +  CaCL  =  CaC204  +  2NH.C1 . 

n.  Application  of  the  above  reactions  to  the  Grouping  of  the  Metals  for 
Analysis. 

o.  A  study  of  the  limit  of  visible  precipitation  with  several  reagents 
upon  a  particular  metal,  or  upon  a  number  of  metals. 

p.  A  study  of  the  analysis  of  the  individual  metals  and  acids;  combining 
them,  and  effecting  their  separation  and  detection.  The  new  work  should 
be  followed  by  the  analysis  of  "unknown5'  mixtures  prepared  by  the 
teacher,  to  illustrate  the  new  work  and  to  give  an  instructive  review  of  the 
preceding  work.  The  order  of  the  study  of  the  metals  and  acids  may  be 
varied  greatly.  In  no  case  should  the  metals  of  a  whole  group  be  studied 
without  considering  the  relations  to  the  other  groups. 

q.  The  study  in  the  class  room  of  Oxidation  and  Reduction,  with  work 
in  the  laboratory  to  illustrate. 

r.  The  study  of  problems  in  Synthesis  involving  analytical  separations, 
accompanied  by  laboratory  experiments. 

s.  The  analysis  of  a  series  of  Dry  "  Unknown  "  Mixtures. 

t.  A  special  study  of  the  analysis  of  Phosphates,  Oxalates,  Eorates, 
Silicates,  etc.,  and  certain  of  the  Rarer  Metals. 

u.  The  analysis  of  mixtures  in  solution,  illustrating  Oxidation  and 
Reduction, 


PART  II.-THE  METALS. 


THE  SILVER  AND  TIN  AND  COPPER  GROUPS. 

(FIRST  AND  SECOND  GROUPS.) 

§47.  The  Silver  group  (first  group)  includes  the  metals  whose  chlorides 
are  insoluble  in  water  and  which  are  precipitated  from  solutions  upon  tlK 
addition  of  hydrochloric  acid  or  soluble  chlorides :  Pb,  Hg',  Ag . 

The  Tin  and  Copper  group  (second  group)  includes  those  metals  whos« 
sulphides  are  precipitated  by  hydrosulphuric  acid  from  solutions  acid  witl 
dilute  hydrochloric  acid,  and  whose  chlorides  (soluble  in  water  for  the 
most  part)  are  not  precipitated  by  hydrochloric  acid  or  soluble  chlorides. 

Lead*  Pb  207.20  Germanium  Ge  72.5 

Mercury  Hg  200.6  Iridium  Ir  193.1 

Silver  Ag>  107.88  Osmium  Os  190.9 

Arsenic  As  74.96  Palladium  Pd  106.7 

.Antimony  Sb  120.2  Ehodium  Bh  102.9 

Tin  Sn  H8.7  Kuthenium  Bu  101.7 

Gold  Au  197.2  Selenium  Se  79.2 

Platinum  Pt  ?95.2  Tellurium  Te  127.5 

Molybdenum  Mo  96.0  Tungsten  "W  184.0 

Bismuth  Bi  208.0  Vanadium  V  51.0 

Copper  Cu  63.57 

Cadmium  Cd  U2.4 

§48.  Owing  to  the  partial  solubility  of  lead  chloride  in  water,  it  is  nevei 
completely  precipitated  in  the  first  group;  hence  it  must  also  be  tested 
for  in  the  second  group.  Monovalent  mercury  belongs  to  the  first  group 
and  divalent  mercury  to  the  second.  Silver,  then,  is  the  only  exclusively 
first-group  metal. 

§49.  The  metals  included  in  these  groups  are  less  strongly  electro- 
positive than  those  of  the  other  groups.  Only  bismuth,  antimony,  tin. 
and  molybdenum  decompose  water,  and  these  only  slowly  and  at  high 
temperatures.  The  oxides  of  silver,  mercury,  gold,  platinum,  and  palla- 
dium are  decomposed  below  a  red  heat.  Copper,  lead,  and  tin  tarnish  by 

*  In  this  list  of  the  metals  of  the  Silver,  Tin  and  Copper  Groups  the  more  common,  those  in 
the  first  column,  are  arranged  in  the  order  of  their  discussion  and  separation  in  analysis.  The 
rare  metals  are  arranged  in  alphabetic  order,  but  are  discussed  in  order  of  their  relations  to 
each  other,  beginning  at  §  1O4. 


28  GENERAL  DISCUSSION.  §50. 

oxidation  in  the  air.  In  general,  these  metals  do  not  dissolve  in  acids 
with  evolution  of  hydrogen,  or  do  so  with  difficulty.  Nitric  acid  is  the 
best  solvent  for  all,  except  for  antimony  and  tin,  which  are  rapidly  oxidized 
by  it.  Antimony  may  be  dissolved  by  treatment  with  a  little  strong  nitric 
acid  and  tartaric  acid.  The  best  solvent  for  tin  is  hot  strong  hydrochloric 
acid.  Concerning  the  separation  and  detection  of  the  metals  of  these 
groups  by  electrolysis,  see  Schmucker,  Z.  anorg.,  1891,  5,  199,  and  Cohen, 
J.  Soc.  Ind.,  1891,  10,  327  (§12). 

§50.  Mercury,  arsenic,  antimony,  and  tin  form,  each  two  stable  classes 
of  salts.  Therefore,  the  lower  oxides,  chlorides,  etc.,  of  these  metals  act 
as  reducing  agents;  and  their  higher  oxides,  chlorides,  etc.,  as  oxidizing 
agents,  each  to  the  extent  of  its  chemical  force.  Arsenic,  antimony,  tin, 
molybdenum,  and  several  of  the  rare  metals  of  these  groups  enter  into 
acid  radicles,  which  form  stable  salts.  Arsenic,  selenium  and  tellurium 
are  metalloids  rather  than  metals.  Arsenic,  antimony,  and  bismuth  belong 
to  the  Nitrogen  Series  of  Elements. 

§51.  A  large  proportion  of  the  compounds  of  these  metals  are  insoluble 
in  water.  Of  the  oxides  or  hydroxides,  only  the  acids  of  arsenic  are 
soluble  in  water.  The  only  insoluble  chlorides,  bromides,  and  iodides  are 
in  these  groups.  The  sulphides,  carbonates,  oxalates,  phosphates,  borates, 
and  cyanogen  compounds  are  insoluble.  Most  of  the  so-called  soluble 
compounds  of  bismuth,  antimony,  and  tin,  and  some  of  those  of  mercury, 
dissolve  only  in  acidulated  water,  being  decomposed  by  pure  water,  with 
formation  of  insoluble  basic  salts. 

§52.  Among  the  many  soluble  double  salts  of  the  metals  of  these  groups 
are  especially  to  be  mentioned  the  double  iodides  with  KI  and  the  iodides 
of  Pb  ,  Hg1 ,  Ag ,  Bi  and  Cd  .  Platinum  forms  a  large  number  of  stable 
double  chlorides,  soluble  and  insoluble;  antl  gold  forms  double  chlorides, 
cyanides,  etc. 

§53.  The  oxides  of  arsenic  act  as  acid  anhydrides  and  form  soluble  salts 
with  the  alkalis;  oxides  of  antimony,  tin,  and  lead,  are  soluble  in  the  fixed 
alkalis ;  oxides  of  silver,  copper,  and  cadmium,  in  ammonium  hydroxide.  Me- 
tallic lead,  like  zinc,  dissolves  in  the  fixed  alkalis  with  evolution  of  hydrogen. 

§54.  The  solubility  of  certain  sulphides  in  the  alkali  sulphides  forming 
sulpho  salts  or  double  sulphides,  separates  tho  metals  of  the  second  group 
into  two  divisions.  A  (copper  group) — Kg,  Pb,  Bi,  Cu,  Cd,  Os,  Pd,  Rh, 
and  Ru;  sulphides  not  soluble  in  yellow  ammonium  sulphide;  and  B 
(tin  group)— As,  Sb,  Sn,  Ge,  An,  Ir,  Mo,  Pt,  Se,  Te,  W,  and  V;  sulphides 
soluble  in  yellow  ammonium  sulphide. 

§55.  Mercury,  antimony,  silver,  and  gold  do  not  form  hydroxides.  The 
oxides  of  gold  a.re  very  unstable. 

§56.  The  metals  of  these  groups  are  all  easily  reduced  to  the  metallic 
state  by  ignition  on  charcoal.  Except  mercury  and  arsenic,  which  vaporize 


§57,  4.  LEAD.  29 

readily,  and  certain  rarer  metals  difficultly  fusible,  the  reduced  metals  melt 
to  metallic  grains  on  the  charcoal. 

THE  SILVER  GKOUP  (FIRST  GROUP). 

Lead,  Mercury  (Mercurosiun),  Silver. 

§57.  Lead   (Plumbum  Pb  =  207.20.     Valence  two  and  foul. 

1.  Properties.—  Specific  gravity,   11.34  (Reich,  J.  pr.,   1859,78,  328).     Melting 
point,  327.4°.     B.   S.,   circular  No.  35,  2nd  ed.,    1915.     Vaporization  is  said  to 
take  place  at  360°  (Demarcay,  C.  r.,  1882,  95,  183).     Boiling  point  about  1525° 
(H.  C.  Greenwood,  C.  N.,  39,  49).     It  can  be  distilled   in   vacuo,  (Schuller,  B., 
1883,  16,  1312). 

Pure  lead  is  almost  white,  soft,  malleable,  very  slightly  ductile;  freshly  cut 
surfaces  tarnish  in  the  air  from  formation  of  a  film  of  oxide.  The  presence  of 
traces  of  most  of  the  other  metals  makes  lead  sensibly  harder.  It  is  a  poor  con- 
ductor of  heat  and  electricity,  and  forms  alloys  with  most  metals;  lead  and  tin  in 
various  proportions  form  solder  and  pewter;  lead  and  arsenic  form  shot  metal; 
lead  and  antimony,  type  metal;  lead,  bismuth,  tin  $  nd  cadmium  form  easily 
fusible  alloys  of  low  melting  points  (minimum  55.5°;  Ch.  Z.,  30,  1139-1143; 
J.  Soc.  Ind.,  25,  1221).;  bell  metal  consists  of  tin,  copper,  lead  and  zinc. 

2.  Occurrence.— It  is  rarely  found  native  (Chapman,  Phil.  Mag.,  1886,  (4),  31, 
176);    the  most  abundant  lead  mineral  is  galena,  PbS;    it  also  occurs  as  cerussite, 
PbCO3;    anglesite,  PbSO4;  pyrqmorphite,  Pb6Cl(Pp4)5;    crocoite,    PbCrO4;    and  in 
many  other  minerals  in   combination  with   arsenic,  antimony,  etc.      The  United 
States  produces  more  lead  than  any  other  country.     Spain  produces  about  one- 
fourth  of  the  world's  supply. 

3.  Preparation. — (a)   From   argentiferous   lead   ores;     after  roasting,  if  neces- 
sary, the  ore  is  smelted  in  a  rectagular  bl  st  furnace  with  a  properly  propor- 
tioned mxiture  of  coke  and  limestone.     The  lead  (base  bullion)  produced  is  de- 
silverized and  refined  by  the  Betts  (electrolytic)    process,  or  desilverized  by  the 
Parkes  process  and  subsequently  refined  in  a  reverberatory  furnace.     (6)   From 
galena,  by  the  roast-reaction  process  in  reverberatory  furnaces,  and  ore  hearths; 
the  ore  is  roasted  with  access  of  air,  forming  variable  quantities  of  PbSO4,  PbO, 
and  PbS.     Air  is  then  excluded  and  the  temperature  raised,  the  sulphur  of  the 
sulphide  then  reduces  both  oxide  and  sulphate  with  formation  of  SO2^ 

PbSO4+    PbO  +  TPbS  -  5Pb  +  3SO2 
In  variations  of  this  process  carb  n  is  used  to  aid  in  the  reduction. 

4.  Oxides.— Lead  forms  four  oxides,  Pb2O  ,  PbO  ,   Pb02  ,   and  Pb3O4  .     Lead 
suboxlde  (Pb2O)  is  little  known:  it  is  the  black  powder  formed  when  PbC2O4  is 
heated  to  300°,  air  being-  excluded.     Lead  oxide  (litharge,  or  massicot)  is  formed 
by  intensely  igniting-  in  the  air  Pb  ,  Pb2O  ,  PbO2  ,  Pb3O4  ,  Pb(OH)2  ,  PbCOri  , 
PbC,04  ,  or'Pb(NO:!)2  •     It  has  a  yellowish-white  color,  melts  at  a  red  heat,  and 
is  volatile  at  a  white  heat. 

Trilead  tetroxide   (red  lead  or  minium),  Pb3O4  ,  is   formed  by  heating1  PbO 
to  a  dull-red  heat  with  full  access  of  air  for  several  hours.     Strong,  non-reduc- 
ing- acids,  such  as  HNO3  ,  H,SO4  ,  HC1O3  ,  etc.,  convert  it  into  a  lead  salt  and 
Pb02  (a).     But  concentrated  hot  ELSO4  converts  the  whole  into  PbSO4  ,  oxygen 
being-   evolved    (&).     But  with   the   dilute   acid   and   reducing   agents,    such    as 
C3H5(OH)3,  CGH12O,,  ,  H2C204  ,  H2C4H4O6  ,   Zn  ,   Al  ,   Cd  ,  Mg  ,   As,   Pb  ,   etc., 
it  is  all  reduced  to  the  dyad  lead  without  evolution  of  oxygen  (c),  (d),  and  (e). 
Hydracids  usualty  reduce  the  lead  and  are  themselves  oxidized  (f). 
(a)      Pb304  +  2H2S04  (dilute)  =  PbO2  +  2PbS04  +  2H20 
(5)      2Pb304  -f  6H2SO4  (concentrated  and  hot)  =  6PbS04  +  6H2O  +  03 

(c)  Pb3O4  +  H2C204  +  6HN03  =  3Pb(NO8)a  +  4H20  +  2C02 

(d)  10Pb304  +  As4  +  30H,S04  =  30PbS04  +  4H3AsO4  +  24H20 

(e)  Pb304  +  Zn  +  4H2SO4  =  3PbSO4  +  ZnSO4  +  4H20 

(f)  Pb304  +  8HC1  =  3PbCl8  +  Cl,  +  4H20 

The  valence  of  Pb3O4  is  best  explained  by  the  theory  that  it  is  a  union  of  the 
dyad  and  tetrad  (Pb"  and  Pbrv)  ,  Pb8O,  =  3PbO  + 


30  LEAD.  §57, 5a. 

Lead  dioxide  or  peroride,  PbO,  ,  is  formed:  (1)  by  fusion  of  PbO  with  KC1O, 
or  KNO,  ;  (2)  by  fusing  Pb304  with  KOH  :  (3)  by  treating-  any  compound  of 
Pb"  with  Cl ,  Br  ,  K3Fe(CN)6  ,  KMn04  ,  or  H202  in  presence  of  KOH;  (4)  by 
treating  Pb304  with  non-reducing  acids: 

Pb304  +  4HN03  =  PbO2  +  2Pb(N03)2  +  2H2O. 

Ignition  forms  first  Pb304  and  above  a  red  heat  PbO,  oxygen  being  given  oft'. 
It  dissolves  in  acids  on  same  conditions  as  Pb304  .  Very  strong  solution  of 
potassium  hydroxide,  in  large  excess,  dissolves  it,  with  formation  of  "  potassium 
plumbate,"  K2Pb03  .  Lead  dioxide  is  a  powerful  oxidizing  agent,  one  of  the 
strongest  known.  Digested  with  ammonium  hydroxide,  it  forms  lead  nitrate 
and  water.  Triturated  with  one-sixth  of  sulphur,  or  tartaric  acid,  or  sugar, 
it  takes  fire;  with  phosphorus,  it  detonates. 

5.  Solubilities. — a. — Meto1* — Nitric  acid  is  the  proper  solvent  for  metallic  lead, 
the  lead  nitrate  formed  is  readily  soluble  in  water  but  insoluble  in  concentrated 
nitric  acid  *;  hence  if  the  concentrated  acid  be  used  to  dissolve  the  lead,  a 
wrhite  residue  of  lead  nitrate  will  be  left  which  dissolves  on  the  addition  ef 
water.  If  concentrated  and  hot,  the  nitric  acid  is  reduced  to  NO  which,  on  contact 
with  the  oxygen  of  the  air,  becomes  N2Oa  (§241,  6).  The  reactions  are  as  follows: 

3Pb  +  8HNO3  =  3Pb(NO3)2+2NO  +  4H2O  4NO  +  O2  =  N2O3 

Dilute  sulphuric  acid  is  without  action,  the  concentrated  acid  is  almost  without 
action  in  the  cold  (Calvert  and  Johnson,  J.  C.,  1863,  16,  66),  but  the  hot  concen- 
trated acid  slowly  changes  the  metal  to  the  sulphate  with  evolution  of  sulphur 
dioxide,  a  portion  of  the  salt  being  dissolved  in  the  acid,  precipitating  on  the 
addition  of  water.  The  following  reaction  takes  place  (§266,  6 A): 

Pb+2H2SO4  =  PbSO4  +  SO2  +  2H2O. 

Hydrochloric  acid  very  slowly  dissolves  the  metal  (more  rapidly  when  warmed),, 
evolving  hydrogen;  the  chloride  formed  dissolves  in  the  acid  in  quantities  depend- 
ing upon  conditions  of  temperature  and  concentration  (c) .  The  halogens  readily 
attack  the  metal  forming  the  corresponding  haloid  salts.  Alloys  of  lead  are  best 
dissolved  by  first  treating  with  nitric  acid;  if  a  white  residue  is  left  it  is  washed 
with  water  and,  if  not  dissolved,  it  is  then  treated  with  hydrochloric  acid,  in 
which  it  will  usually  be  soluble. 

Water  used  for  drinking  or  cooking-  purposes  should  not  be  allowed  to  stand 
in  lead  pipes.  Pure  water  free  from  air  is  Avithout  action  upon  pure  lead,  but 
water  containing  air  and  carbon  dioxide  very  slowly  attacks  lead,  forming  the 
hydroxide  and  basic  carbonate.  This  action  is  promoted  by  the  presence  of 
salts,  as  ammonium  nitrate,  nitrite,  chloride,  etc.;  the  action  seems  to  be 
hindered  by  the  presence  of  sulphates. 

6. — Oxides. — Lead  oxide,  I i Marge,  PbO  ,  and  the  hydroxides,  2PbO.H2O; 
3PbO.H2O,  are  readily  dissolved  or  transposed  by  acids  forming  the  correspond- 
ing salts,  i.  €.,  PbO  +  H2S04  =  PbS04  -j-  H.,O  ."  The  oxide  and  hydroxide  are 
soluble  in  about  7000  parts  of  water,  to  which  they  impart  arr  alkaline  reaction. 
They  are  soluble  in  the  fixed  alkalis  forming  plumbites;  soluble  in  certain  salts 
as  NH4C1  ,  CaClo  ,  and  SrCL  (Andre,  C.  r.,  1883,  96,  435;  1887,  104,  359);  very 
soluble  in  lead  acetate,  forming  a  strongly  alkaline  solution  of  basic  lead  acetate. 

Lead  dioxide,  Pb02  ,  lead  pero-ride,  is  insoluble  in  water  or  nitric  acid:  it  is 
dissolved  by  the  halogen  hydracids  with  liberation  of  the  halogen  and  reduction 
of  the  lead  forming  a  dyad  salt:  PbO,  +  4HC1  =  PbCL  +  C12  +  2ELO;  it  is 
attacked  by  hot  concentrated  sulphuric  acid,  forming  the  sulphate  and  liberat- 
ing oxygen;  it  is  soluble  in  glacial  acetic  acid  forming  Pb(C2H302)4  ,  unstable 
(Hutchinson  and  Pollard,  J.  C.,  1896,  69,  212).  Some  of  the  salts  of  the  tetrad 
lead  seem  to  be  .formed  when  the  peroxide  is  treated  with  certain  acids  in  the 
cold.  They  are,  however,  very  unstable,  being  decomposed  to  the  dyad  salt 
upon  warming  (Fischer,  J.  C.,  1879,  35,  282;  Nickels,  A.  Ch.,  1867,  (4),  10,  328). 
The  peroxide  is  slowly  soluble  in  the  fixed  alkali  hydroxides  forming  plum- 
bates,  i.  e.,  PbO2  +  2KOH  =  K2PbO3  +  H,O  . 

Trilead  tetroxide,  Pb304  ,  red  lead,  nihunm,  is  insoluble  in  water,  is  at- 
tacked by  nearly  all  acids  in  the  cold  forming  the  corresponding  dyad  lead 

*  The  solubility  of  a  salt  is  lessened  by  the  presence  of  another  substance  having  an  ion  in 
common  with  it  (§  45).  In  some  cases,  as  with  Pbla  and  KI,  this  Is  offset  In  concentrated 
solution  by  the  formation  of  a  complex  compound. 


§57,  5a.  LEAD.  31 

salt  and  ead  peroxide,  PbO2 .  Upon  further  treatment  with  the  acids  using 
heat  the  lead  peroxide  is  decomposed  as  described  above.  The  presence  of 
many  reducing  agents,  as  alcohol,  oxalic  acid,  hydrogen  peroxide,  etc.,  greatly 
facilitates  the  solution  of  red  lead  or  lead  peroxide  in  acids,  i.  e.,  nitric  acid 
does  not  dissolve  lead  peroxide,  but  if  a  few  drops  of  alcohol  be  added  the 
solution  is  readily  obtained  upon  warming,  the  lead  being  reduced  and  then 
converted  into  the  soluble  nitrate. 

c. — Salts. — The  carbonate,  borate,  cyanide,  ferrocyanide,  phosphate,  sul- 
phide, sulphite,  iodate,  chromate,  and  tannate  are  insoluble  in  water. 
The  sulphate  is  soluble  in  about  21,000  parts  of  water  at  18°  ( Kohlrausch 
and  Rose,  Z.  phys.  Ch.f  1893,  12,  241),  the  presence  of  HN03  or  HC1  in- 
creasing its  solubility  in  water;  it  is  insoluble  in  alcohol  even  when  quite 
dilute;  sparingly  soluble  in  concentrated  H2S04 ,  from  which  solution  it  i  » 
precipitated  by  the  addition  of  water  or  alcohol;  less  soluble  in  dilute  11,80^ 
than  in  water;  soluble  in  682  parts  10  per  cent  HC1 ,  in  35  parts  31.5  per 
cent  (Eodwcll,  J.  C.,  1862,  15,  59);  transposed  and  dissolved  by  excess  of 
HC1 ,  HBr,  or  HI  forming  the  corresponding  haloid  salt;  insoluble  in 
HF  (Ditte,  A.  (77?,.,  1878,  (5),  14,  190);  soluble  in  ammonium  sulphate, 
nitrate,  acetate,  tartrate  and  citrate,  and  from  these  solutions  not  readily 
precipitated  by  ammonium  hydroxide  or  sulphate  (Fleischer,  J.  C.,  1876, 
29,  190;  Woehler,  A.,  1840,  34,  235).  The  sulphate  is  almost  completely 
transposed  to  the  nitrate  by  standing  several  days  with  cold  concentrated 
nitric  acid  (Rodwell,  I.  c.).  The  oxalate  is  sparingly  soluble  in  water,  insol- 
uble in  alcohol.  The  ferricyanide  is  very  slightly  soluble  in  cold  water,  more 
soluble  in  hot  water.  The  chloride  is  soluble  in  85  parts  water  at  20°  and  in 
32  parts  at  80°  (Ditte,  C.  r.}  1881,  92,  718).  The  bromide  is  soluble  in  166 
parts  water  at  10°,  in  about  45  parts  at  80°.  The  iodide  is  soluble  in  1235 
parts  water  at  ordinary  temperature,  and  in  194  parts  at  100°  "(Denot,  J . 
;//•.,  1834,  1,  425).  The  chloride  is  less  soluble  in  dilute  HC1  or  H.SO,  than 
in  water,  but  is  more  soluble  in  the  concentrated  acids  (Ditte,  I.  c.) ;  HNO:>> 
increases  the  solubility  of  the  chloride  more  and  more  as  the  HNO.{  is 
stronger.  The  chloride  is  less  soluble  in  a  solution  of  NaCl  than  in  water 
(Field,  J.  0.,  1873,  26,  575);  soluble  in  NH4C1  —90  grams  dissolving  in  200 
grams  NH4C1  with  200  cc.  water  (Andre,  (7.  r.,  1893,  96,  435).  The  chloride, 
bromide,  and  iodide  are  insoluble  in  alcohol.  The  iodide  is  moderately 
soluble  in  solutions  of  alkali  iodides;  it  is  decomposed  by  ether.  The 
basic  acetates  are  permanently  soluble  if  carbonic  acid  is  strictly  excluded. 
The  basic  nitrates  are  but  slightly  soluble  in  water,  and  are  precipitated 
on  adding  solutions  of  KNOn  to  a  solution  of  basic  lead  acetate. 

The  relative  insolubility  of  PbCl2  in  cold  water  or  in  dilute  HC1  makes 
it  possible  to  precipitate  the  most  of  the  lead  (by  means  of  HC1)  from 
solutions  containing  also  the  other  metals  of  the  Silver  Group;  while  its 
solubility  in  hot  water  is  the  means  of  its  separation  from  the  other 
chlorides  of  that  group  (§61).  The  lead  is  separated  and  identified  in 
the  second  group  as  the  insoluble  sulphate.  (§95). 


3%  LEAD.  §57, 6. 

6.  Reactions,  a. — Fixed  alkali  hydroxides  precipitate,  from  solutions  of 
lead  salts,  basic  lead  hydroxide  (1),  Pb,0(OH)2  (Schaffner,  A.,  1844,  51,  175), 
white,  soluble  *  in  excess  of  the  reagent  as  plumbite  (2)  (distinction  from 
silver,  mercury,  bismuth,  copper,  and  cadmium).  The  normal  lead  hy- 
droxide, Pb(OH)2 ,  may  be  formed  by  adding  a  solution  of  a  lead  salt  to 
a  solution  of  a  fixed  alkali  hydroxide. 

(1)  2Pb(N03)2  +  4KOH  =  Pb,0(OH)2  +  4KN03  +  H20 

(2)  PbaO(OH),  +  4KOH  =  2K,Pb02  +  3H20  . 

Ammonium  hydroxide  precipitates  white  basic  salts,  insoluble  in  water 
and  in  excess  of  the  reagent  (distinction  from  silver,  copper,  and  cad- 
mium); with  the  chloride  the  precipitate,  insoluble  in  water,  is 
PbClo.PbO.HL>0  (Wood  and  Bordeu,  C.  N.,  1885,  52,  43);  with  the  nitrate 
2PbO.Pb(N03)2  (D.,  2,  2,  558).  With  the  acetate,  in  solutions  of  ordinary 
strength,  excess  of  ammonium  hydroxide  (free  from  carbonate)  gives  no 
precipitate,  the  soluble  tribasic  acetate  being  formed. 

Alkali  carbonates  precipitate  basic  lead  carbonate,  white,  the  composition 
varying  with  the  conditions  of  precipitation.  With  excess  of  the  reagent 
and  in  hot  concentrated  solutions  the  precipitate  consists  chiefly  of 
Pb;j(OH)o(CO.,)2  .  Precipitation  in  the  cold  approaches  more  nearly  to  the 
normal  carbonate  (Lefort,  Pharm.  J.,  1885,  (3),  15,  26).  Solutions  of  lead 
salts  when  boiled  with  freshly  precipitated  barium  carbonate  are  com- 
pletely precipitated.  Carbon  dioxide  precipitates  the  basic  acetate  but 
not  completely. 

b. — Oxalic  acid  and  alkali  oxalates  precipitate  lead  oxalate,  PbC204,  white, 
from  solutions  of  lead  salts,  soluble  in  nitric  acid,  insoluble  in  acetic  acid. 
A  solution  of  lead  acetate  precipitates  a  large  number — and  a  solution  of 
lead  subacetate  a  still  larger  number — of  organic  acids,  color  substances, 
resins,  gums,  and  neutral 'principles.  Indeed  it  is  a  rule.,  with  few  excep- 
tions, that  lead  subacetate  removes  organic  acids  (not  formic,  acetic, 
butyric,  valeric,  or  lactic).  Tannic  acid  precipitates  solutions  of  lead 
acetate,  and  of  the  nitrate  incompletely,  as  yellow-gray  lead  tannate, 
soluble  in  acids. 

Soluble  cyanides  precipitate  lead  cyanide,  Pb(CN),  ,  white,  sparingly  soluble 
hi  a  large  excess  of  the  reagent  and  reprecipitated  on  boiling.  Potassium  ferro- 
cvanide  precipitates  lead  ferrocyanide,  Pb2Fe(CN)6  ,  white,  insoluble  in  water 
or  dilute  acids.  Potassium  ferricyanide  precipitates  from  solutions,  not  too 
dilute,  lead  ferricyanide,  Pb3(Fe(CN)6)2,  white,  sparingly  soluble  in  water,  soluble 
in  nitric  acid.  Solutions  of  lead  salts  are  precipitated  by  potassium  sulpho- 
cyanate  as  lead  sulphocyanate,  Pb(CNS)2,  white,  soluble  in  excess  of  the  reagent 
and  in  nitric  acid. 

c. — Lead  nitrate  is  very  soluble  in  water,  the  solution  dissolving  the  oxide  to 
form  a  basic  nitrate,  which  may  also  be  formed  by  precipitating  lead  acetate  with 

*  Nearly  all  the  salts  are  soluble  in  the  fixed  alkali  hydroxides,  PbS  forming  almost  the  only* 
notable  exception. 


§57,  Ge.  LEAD.  33 

potassium  nitrate.  The  solubility  of  lead  nitrate  is  greatly  increased  by  the 
presence  of  the  nitrates  of  the  alkalis  and  of  the  alkaline  earths,  a  complex 
compound  beiny  formed  (  Le  I '.lane  and  A'oyes,  Z.  pliy*.  ('//.,  1890,  6,  385). 

</. — The  higher  oxides  of  lead  are  all  reduced  by  hypOpUosphOTOUS  acid,  lead 
phosphate  being  formed.  Lead  phosphite,  PbHP03  ,  white,  is  formed  by 
neai-U  neutralizing  phosphorous  acid  with  lead  carbonate  or  precipitating 
Na,HPO,  \\ith  Pb(N03),  (Ainat,  ('.  r.,  1890,  110,  901).  Sodium  phosphate, 
Na.HPO,  ,  precipitates  from  solutions  of  lead  acetate  the  tribasic  lead  phosphate, 
Pb3(PO4),  ,  white,  insoluble  in  the  acetic  acid  which  is  set  free  (D.,  2,  2,  562): 
3Pb(C?HaOa)a  +  I^Na.HPO,  =  Pb3(P04)2  +  4NaC2H302  +  2HC2H,02.  The  same 
precipitate  is  formed  wh^ii  sodium  phosphate  is  added  to  lead  nitrate,  soluble 
in  nitric  acid,  insoluble  in  acetic  acid.  Lead  phosphate  is  also  precipitated 
upon  the  addition  of  phosphoric  acid  to  solutions  of  lead  acetate  or  lead  nitrate. 
The  pyrophosphatc,  Pb,P,07  ,  white,  amorphous,  is  formed  by  precipitating  a 
lead  solution  with  Na,P,07  ,  soluble  in  excess  of  the  precipitant,  in  nitric  acid, 
and  in  potassium  hydroxide;  insoluble  in  ammonium  hydroxide  and  in  acetic 
acid  (Gerhardt,  A.  (.'/».,  1849,  (3),  25,  305).  The  metaphosphate,  Pb(PO3)2  , 
white,  crystalline,  is  obtained  by  the  action  of  NaPO3  upon  Pb(N03)2  in  excess. 

e. — Hydrosulphuric  acid  and  the  soluble  sulphides  precipitate — from 
neutral,  acid,  or  alkaline  solutions  of  lead  salts — had  sulphide,  PbS, 
brownish  black,  insoluble  in  dilute  acids,  in  alkali  hydroxides,  carbonates, 
or  sulphides.  Freshly  precipitated  CdS,  MnS,  FeS,  CoS ,  and  NiS  also 
give  the  same  precipitate.  Hydrosulphuric  acid  and  the  soluble  sulphides 
transpose  all  freshly  precipitated  lead  salts  to  lead  sulphide.*  Moder- 
ately dilute  nitric  acid— 15  to  20  per  cent— dissolves  lead  sulphide  with 
separation  of  sulphur  (1),  some  of  the  sulphur,  especially  if  the  nitric  acid 
be  concentrated,  is  oxidized  to  sulphuric  acid,  which  precipitates  a  portion 
of  the  lead  (2),  unless  the  nitric  acid  be  sufficiently  concentrated  to  hold 
that  amount  of  lead  sulphate  in  solution.  The  oxidation  of  sulphur  always 
occurs  when  nitric  acid  acts  upon  sulphides,  and  in  degree  proportional 
to  the  strength  of  acid,  temperature,  and  duration  of  contact. 

(?)     6PbS  +  16HN03  =  6Pb(N03)2  -f  3S2  +  4NO  +  8H2O 
(2)     3PbS  +  SHN03  =  3PbS04  -f  8NO  +  4H2O 

In  solutions  too  strongly  acidulated,  especially  with  hydrochloric  acid, 
either  no  precipitation  takes  place,  or  a  brick-red  double  salt,  Pb2SCl2 , 

*  The  condition  for  equilibrium  is  that  a  certain  ratio  of  concentration  exist  between  the  ions, 
in  the  case  of  PbSO4  between  the  S  ions  and  the  SO4  ions.  These  concentrations  are  the  same 
as  those  in  a  solution  obtained  by  digesting  the  two  salts,  PbSO4  and  PbS.  together  in  water. 
PbSO4  dissolves  more  freely  than  PbS.  and  for  equilibrium  therefore  C~Q  .,  must  be  corres- 
pondingly greater  than  cs->.  But  adding  H2S  or  a  soluble  sulphide  to  PbS<>4  gives  just  the 
opposite  of  this  condition,  and  transformation  accordingly  results,  increasing  the  SO4"  con- 
ce  itrati'Mi  by  formation  of  soluble  sulphate  and  decreasing  the  S"  concentration  by  precipita- 
tion of  PbS,  until  the  equilibrium-ratio  is  produced  or,  if  the  quantity  of  Pl>SO4  present  is  in- 
s'lTu-ient  for  this,  until  all  the  PbSO4  has  been  transformed  to  sulphide.  On  tlr>  other  hand, 
treatment  of  PbS  with  a  very  large  excess  of  H2SO4  will  cause  the  reverse  action,  S  ions  going 
into  solution  until  the  same  equilibrium  results  as  before. 

The  general  principle  is  then  that  unless  a  constituent  of  the  more  soluble  substance  is  in  great 
preponderance  in  the  solution  the  least  soluble  of  two  or  more  possible  products  will  always  be 
formed.     This  principle  determines  the  direction  in  which  a  reaction  takes  place; 
AgCl  +  KI=  Agl  +  KCl  ;  CaSO4  +  Na2COi  =  CaCOa  + Na2SO4  (U4). 


34  LEAD.  §57,  6/. 

is  formed,  the  precipitation  being  incomplete.  In  neutral  solutions  con- 
taining 100,000  parts  of  water  lead  is  revealed  as  the  sulphide;  a  test 
which  is  much  more  delicate  than  the  formation  of  the  sulphate. 

Ferric  chloride  decomposes  lead  sulphide,  forming-  lead  chloride,  ferrous 
chloride  and  sulphur.  The  reaction  takes  place  in  the  cold  and  rapidly  when 
warmed  (Gabba,  C.  (7.,  1889,  667). 

When  galena,  PbS  ,  is  pulverized  with  fused  KHSO4  ,  H,S  is  evolved  (Jan- 
nettaz,  J.  C.,  1874,  27,  188). 

Lert<l  thiosulphate,  PbS,O3  ,  white,  is  precipitated  by  adding-  sodium  thiosul- 
phate  to  solutions  of  lead  salts;  the  precipitate  is  readily  dissolved  in  an  excess 
of  the  reag-ent,  forming-  the  double  salt,  PbSoO^NaXo.,  (Lenz,  A.,  1841,  40, 
94) ;  on  boiling-,  all  the  lead  is  slowly  precipitated  as  sulphide  (Vohl,  A.,  1855, 
96,  237). 

Sodium  sulphite  precipitates  lead  sulphite,  PbS03 ,  white,  less  soluble  in 
water  than  the  sulphate,  slightly  soluble  in  sulphurous  acid;  decomposed 
by  sulphuric,  nitric,  hydrochloric,  and  hydrosulphuric  acids  and  by  alkali 
sulphides;  not  decomposed  by  cold  phosphoric  and  acetic  acids. 

Sulphuric  acid  and  soluble  sulphates  precipitate  from  neutral  or  acid 
solutions,  lead  sulphate,  PbS04 ,  white,  not  readily  changed  or  permanently 
dissolved  by  acids,  except  hydrosulphuric  acid,  yet  slightly  soluble  in 
strong  acids  (5c).  Soluble  in  the  fixed  alkalis  and  in  most  ammonium 
salts,  especially  the  acetate,  tartrate,  and  citrate  (Woehler,  .1.,  1$40,  34, 
235).  Soluble  in  warm  sodium  thiosulphate  solution,  in  hot  solution 
decomposed,  lead  sulphide,  insoluble  in  thiosulphate,  being  formed  (dis- 
tinction and  separation  from  barium  sulphate,  which  does  not  dissolve  in 
thiosulphates). 

The  test  for  lead  as  a  sulphate  is  from  five  to  ten  times  less  delicate 
than  that  with  hydrosulphuric  acid;  but  lead  is  quantitatively  separated 
as  a  sulphate,  by  precipitation  with  sulphuric  acid  in  the  presence  of 
alcohol,  and  washing  with  alcohol.  When  heated  with  potassium  ehromate 
transposition  takes  place  and  yellow  lead  chromate  is  formed  (//).  Excess 
of  potassium  iodide  transposes  lead  sulphate  (/),  a  distinction  of  load  from 
barium.  Repeated  washing  of  lead  sulphate  with  a  solution  of  sodium 
chloride  completely  transposes  the  lead  to  the  chloride  (Matthey.  J.  f'., 
1879,  36,  124).  See  footnote  on  previous  page. 

f. — Hydrochloric  acid  and  soluble  chlorides  precipitate,  from  solutions 
not  too  dilute,  lead  chloride,  PbCL  ,  white.  This  reaction  constitutes  lend 
a  member  of  the  FIRST  GROUP — as  it  also  is  of  the  second.  Tlio  solu- 
bility of  the  precipitate  is  such  (5c)  that  the  filtrate  obtained  in  the  cold 
giTes  marked  reactions  with  hydrosulphuric  acid,  sulphuric  acid,  chro- 
mates,  etc.;  and  that  it  can  be  quite  accurately  separated  from  silver 
chloride  and  mcrcurous  chloride  by  much  hot  water.  Also,  small  propor- 
tions of  lead  escape  detection  in  the  first  group,  while  its  rrninnil  is 
necessarily  accomplished  in  the  second  group. 


Hydrobromic  acid  and  soluble  bromides  precipitate  lend  hronudr,  PbBr2 
white,  somewhat  less  soluble  in  water  than  the  chloride  (5c);  soluble  in 
excess  of  concentrated  potassium  bromide,  as  2KBr.PbBr2 ,  which  is  decom- 
posed and  PbBr2  precipitated  by  dilution  with  water. 

Hydriodic  acid  and  soluble  iodides  precipitate  lead  iodide,  PbI2 ,  bright 
yellow  and  crystalline,  much  less  soluble  in  water  than  the  chloride  or 
bromide  (oc);  soluble  in  hot  moderately  concentrated  nitric  acid  and  in 
solution  of  the  fixed  alkalis;  soluble  in  excess  of  the  alkali  iodides,  by 
forming  double  iodides,  KIPbL  with  small  excess  of  KI ,  and  4KI.PbI2 
will)  greater  excess  of  KI  ;  these  double  iodides  are  decomposed  by  addi- 
tion of  water  with  precipitation  of  the  lead  iodide.  Lead  iodide  is  not 
precipitated  in  presence  of  sodium  citrate;  alkali  acetates  also  hold  it  in 
solution  to  some  extent,  so  that  it  is  less  perfectly  precipitated  from  the 
acetate  than  from  the  nitrate.  Freshly  precipitated  lead  peroxide,  Pb02 . 
gives  free  iodine  when  treated  with  potassium  iodide  (Ditte,  C.  r.,  1881, 
93,  64  and  67). 

In  detecting  lead  as  an  iodide  in  solutions  of  the  chloride  by  precipita- 
liou  with  potassium  iodide  and  recrystallization  of  the  yellow  precipitate 
from  hot  water,  care  must  be  taken  that  the  potassium  iodide  be  not 
added  in  excess  to  form  the  soluble  double  iodides. 

0- — Arsenous  acid  does  not  precipitate  neutral  solutions  of  lead  salts;  from 
alkaline  solutions  or  with  soluble  arsenites  a  bulky  white  precipitate  of  lead 
arsenite  is  formed,  insoluble  in  water,  but  readily  soluble  in  all  acids  and  in  the 
fixed  alkali  hydroxides.  Arsenic  acid  and  soluble  arsenates  precipitate  lead 
ar senate,  white,  from  neutral  or  alkaline  solutions  of  lead  salts,  soluble  in  the 
fixed  alkali  hydroxides  and  in  nitric  acid,  insoluble  in  acetic  acid.  For  the 
composition  of  the  arsenites  and  arsenates  of  lead  see  (D.,  2,  2,  565).  Hot 
potassium  stannite  (SnCl2  in  solution  by  KOH)  gives  with  lead  salts  or  lead 
hydroxide  a  black  precipitate  of  metallic  lead. 

//.—Chromic  acid  and  soluble  chromates— both  K,Cr04  and  K2Cr007— 

24  2          J      7 

precipitate  lead  chromate,  PbCr04 ,  yellow,  soluble  in  the  fixed  alkali 
hydroxides  (distinction  from  bismuth),  insoluble  in  excess  of  chromic  acid 
(distinction  from  barium),  insoluble  in  ammonium  hydroxide  (distinction 
from  silver),  decomposed  by  moderately  concentrated  nitric  and  hydro- 
chloric acids,  insoluble  in  acetic  acid.  The  precipitate  is  formed  as  follows : 

PbCl2  +  K2Cr2O7  +  H2O  =  PbCrO,  +  H2CrO4  +  2KC1 

7.  Ignition. — Insoluble  lead  salts  may  be  tested  by  fusion  in  a  porcelain 
crucible  with  sodium  carbonate.  The  lead  is  converted  into  lead  oxide, 
PbO  (a).  After  fusion  and  digestion  with  warm  water,  the  aqueous  solution 
is  tested  for  acids,  and  the  residue  for  bases  after  dissolving  in  nitric  or 
acetic  acid.  If  charcoal  (or  some  organic  compounds  as  sugar,  tartrates, 
etc.)  be  present,  metallic  lead  is  formed  (b)  ;  and  with  excess  of  charcoal 
the  acid  radicle  may  also  be  changed  (c).  If  the  fusion  with  sodium  carbo- 


36  LEAD.  §57,  8. 

nate  is  made  on  a  piece  of  charcoal,  instead  of  in  a  crucible,  using  the  re- 
ducing flame  of  the  blowpipe,  globules  of  metallic  lead  are  produced  and  at 
the  same  time  the  charcoal  is  covered  with  a  yellow  incrustation  of  lead 
oxide,  PbO. 

(a)     PbCl2  +  Na2C03  =  2NaCl  +  PbO  +  C02 

(6)      2PbS04  +  2Na2C03  +  C  =  2Pb  +  2Na2S04  +  SCO, 

(c)      2PbS04  4-  2Na2C03  +  5C  =  2Pb  +  2Na2S  +  7C02 

/       r          i         v*  \  I  v 

8.  Detection. — Lead  is  precipitated,  incompletely,  from  its  solutions  by 

HC1  as  PbCl2  ;  separated  from  AgCl  and  HgCl  by  hot  water,  and  confirmed 
by  H2S ,  H2S04 ,  K2Cr04 ,  and  KI .  It  is  separated  (in  the  second  group) 
from  As ,  Sb ,  Sn ,  etc.,  by  non-solubility  of  the  sulphide  in  (NH4)2SX  ; 
from  HgS  by  HN03  ;  from  Bi,  Cu,  and  Cd  by  precipitation  with  dilute 
sulphuric  acid.  Insoluble  compounds  are  transposed  by  an  alkali  sulphide, 
being  then  treated  as  lead  in  the  second  group,  or  they  are  examined  by 
ignition  as  described  in  (7). 

9.  Estimation. — (a)  As  an  oxide  into  which  it  is  converted  by  ignition  (if  a 
carbonate  or  nitrate),  or  by  precipitation  and  subsequent  ignition,  (ft)  As  a 
sulphate.  Add  to  the  solution  twice  its  volume  of  alcohol,  precipitate  with 
H2SO4  ,  and  after  washing1  with  alcohol  ignite  and  weigh,  (c)  It  is  converted 
into  an  acetate,  or  sodium  acetate  is  added  to  the  solution,  then  precipitated 
with  K2Cr,07  ,  and  after  drying-  at  100°,  weighed  as  PbCrO4  .  (d)  It  is  con- 
verted into  PbS  ,  free  sulphur  added,  and  after  ignition  in  hydrogen  gas 
weighed  as  PbS  .  (e)  The  lead  is  precipitated  with  standardized  sodium  iodate 
and  the  excess  of  iodate  is  determined  by  retitration.  Lead  iodate  is  less 
soluble  in  water  than  lead  sulphate  (Cameron,  J.  C.,  1879,  36,  484).  (f)  In 
presence  of  bismuth,  ignite  the  halogen  compound,  or  convert  into  a  sulphide 
and  ignite  in  a  current  of  bromine.  The  haloid  salts  of  bismuth  sublime  upon 
ignition  (Steen,  Z.  angew.,  1895,  530).  (g)  Gas  -volumetric  method.  Precipitate  as 
a  chromate,  filter,  wash  and  transfer  to  an  azotometer  with  dilute  sulphuric 
acid  and  estimate  the  amount  of  chromium  by  the  volume  of  oxygen  set  free 
by  hydrogen  peroxide  (Baumann,  Z.  angew.,  1891,  329), 

10.  Oxidation. — Metallic  lead  precipitates  the  free  metals  from  solutions 
of  Hg ,  Ag ,  Au ,  Pt ,  Bi ,  and  Cu .  Lead  as  a  dyad  is  oxidized  to  the 
tetrad  as  stated  in  (4),  also  electrolytically  in  separation  from  Cu  (Nlssen- 
son,  Z.  angew.,  1893,  646).  PbIV  is  reduced  to  Pb°  in  presence  of  dilute 
H2S04  by  nascent  hydrogen,  and  by  all  metals  capable  of  producing  nascent 
hydrogen  (such  as  Al ,  Zn ,  Sn ,  Mg ,  Fe),  and  to  Pb"  by  soluble  compounds 
of  Hg',  Sn",  Sb'",  As'",  (AsH3  gas),  Cu',  Fe",  Cr'"^  Mn",  Mn'",  Mnlv, 
MnVI.  Also  by  H2C204 ,  HN02 ,  H3P02 ,  H3P03 ,  P  ,  S02 ,  H2S ,  HC1 ,  HBr , 
HI ,  HCN" ,  HCNS ,  H4Fe(CN)6 ,  glycerine,  tartaric  acid,  sugar,  urea,  and 
very  many  other  organic  compounds.  In  many  cases  the  reduction  to 
Pb"  or  to  Pb°  takes  place  in  presence  of  KOH .  The  freshly  precipitated 
peroxide  oxidizes  ammonia,  NH3 ,  to  nitrite  and  nitrate  in  the  course  of  a 
few  hours  (Pollacci,  Arch.  Pharm.,  1886,  224,  176). 

From  lead  solutions  Zn ,  Mg ,  Al ,  Co ,  and  Cd  precipitate  metallic  lead. 


§58,  5a.  MERCURY.  37 

§58.  Mercury  (Hydrargyrum)  Hg  =  200.0  .     Valence  one  and  two. 

1.  Properties.— Specific  gravity,  liquid,  13.5953  (Volkmann,  W.  A.,  1881,  13,  209); 
solid,  14.19:52  (Mallet,  /'/'or.  h'.  Nor.,  1877,  26,  71).  Mcltlmj  (freezing)  point,  —38.85° 
(Mallet,  Phil.  Mug.,  1877,  (5),  4,  145).  Boiling  point,  :;.")7.:>,3°  at  760  mm.  (Ramsay 
and  Young1,  •/.  C.,  1885,  47,  657).  It  is  the  only  metal  which  is  a  liquid  at 
ordinary  temperatures,  white  when  pure,  with  a  slightly  bluish  tinge,  and 
having  a  brilliant  silvery  lustre.  The  precipitated  or  finely  divided  mercury 
appears  as  a  dark  gray  powder.  Mercury  may  be  "  extinguished  "  or  "  dead- 
ened," /.  e.,  reduced  to  the  finely  divided  state,  by  shaking  with  sugar,  grease, 
chalk,  turpentine,  ether,  etc.  It  is  slightly  volatile  even  at  —13°  (Regnault, 
C.  r.,  1881,  93,  308);  is  not  oxidized  by  air  or  oxygen  at  ordinary  temperature 
(Shenstone  and  Cundall,  J.  C.,  1887,  51,  619).  The  solid  metal  is  composed  of 
octahedral  and  needle-shaped  crystals,  is  very  ductile  and  is  easily  cut  with  a 
knilV.  Owing  to  its  very  strong  cohesive  property  it  forms  a  convex  surface 
with  glass,  etc.  It  is  a  good  conductor  of  electricity,  and  forms  amalgams  with 
Al  ,  Ba  ,  Bi  ,  Cd  ,  Cs  ,  Ca  ,  Cr  ,  Co  ,  Cu  ,  Au  ,  Fe  ,  Pb  ,  Mg  ,  Mn  ,  Ni  ,  Os  , 
Pd,  Pt,  K,  Ag,  Na,  Tl,  Sn,  and  Zn.  Amalgams  with  special  alloys  of  gold, 
silver,  tin,  and  zinc  are  used  for  filling  teeth. 

2.  Occurrence. — Occasionally  found  native  in  small  globules  associated  with 
cinnabar,  in  the  containing   gangue,    and   as   amalgam    (Ag2Hg3  to  Ag36Hg);  the 
principal  mercury  mineral  is  cinnabar,  HgS.     It  occurs  also  as  calomel,  HgCl , 
generally  associated  with  cinnabar. 

Found  in  Austria,  Spain,  Peru,  China,  Russia,  California,  Texas  and  Oregon. 

3.  Preparation. — The    extraction   of   mercury   from    cinnabar,  which   may  be 
considered  as  practically  the  only  ore  of  this  metal,  is  effected:    (a)  by  oxidation 
with  a  regulated  supply  of  air,  and  volatilization  of  the  liberated  metal,  which 
distils  over  and  is  condensed:    HgS  +  O2  =  Hg  +  SO2 ;    (b)  by    mixing   the  ore 
with    lime,    and    distilling:      4HgS  +  4CaO  =  4Hg  +  3CaS  +  CaSO4 .     (c)  The 
ore  is  heated  with  iron  (smithy  scales):    Hg,  FeS  ,  and  SO2  are  produced.     The 
mercury  is  usually   condensed  in   a   trough  of  water.     Commercial  mercury  is 
freed  from  dirt  and  other  impurities  by  pressing  through  leather  or  by  passing 
through  a  cone  of  writing  or  filter  paper  having  a  small  pin-hole  in  the  apex. 
For  the  separation  of  mercury  from  small  quantities  of  Pb  ,  Sn ,  Zn ,  and  Ag 
without  distilling,  see   Briihl   (B.,   1879,   12,   204),     Meyer   (B.,    1879,    12,    437) 
and  Crafts  (Bl.,  1888,  (2),  49,  856). 

A.  Oxides. — Mercury  forms  two  oxides,  Hg.O  and  HgO  .  Mercurous  oxide, 
Hg20  ,  is  a  black  powder  formed  by  the  action  of  fixed  alkalis  on  mercurous 
salts.  It  is  converted  by  gentle  heat  into  Hg  and  HgO  and  by  a  higher  (red) 
heat,  to  Hg  and  O  .  Mercuric  o,mde,  HgO  ,  is  made  (1)  by  keeping  Hg  at  its 
boiling  point  for  a  month  or  longer  in  a  flask  filled  with  'air;  (2)  by  heating 
HgNO3  or  Hg(N03)2  with  about  an  equal  weight  of  metallic  mercury: 
Hg(N03),  -f  3Hg  =  4 HgO  +  2NO;  (3)  by  precipitating  mercuric  salts  with 
KOH  or  NaOH  .  Made  by  (1)  and  (2)  it  is  red,  by  (3)  yellow.  On  heating  it 
changes  tp  vermillion  red,  then  black,  and  on  cooling  regains  its  original  color. 
A  red  heat  decomposes  it  completely  into  Hg  and  O  .  Mercury  forms  no 
hydroxides. 

5.  Solubilities. — a. — Metal. — Unaffected  by  treatment  with  alkalis.  The  movst 
effective  solvent  of  mercury  is  nitric  acid.  It  dissolves  readily  in  the  dilute 
acid  hot  or  cold;  with  the  strong  acid,  heat  is  soon  generated;  and  with  con- 
siderable quantities  of  material,  the  action  acquires  an  explosive  violence.  At 
ordinary  temperatures,  nitric  acid,  when  applied  in  excess,  produces  normal 
mercuric  nitrate,  but  when  the  mercury  is  in  excess,  and  the  acid  is  cold  and 
dilute,  mercurous  nitrate  is  formed;  in  a1!  cases,  chiefly  nitric  oxide  gas  is  generated. 
Both  mercurous  and  mercuric  nitrates  require  a  little  free  nitric  acid  to  hold  them 
in  solution.  This  free  nitric  acid  gradually  oxidizes  mercurous  to  mercuric, 
making  a  clear  solution  of  Hg(NO3)2  ,  if  there  is  sufficient  HNOa  present,  other- 
wise a  basic  mercuric  nitrate  may  precipitate.  A  solution  of  mercurous  nitrate 
may  be  kept  free  from  mercuric  nitrate  by  placing  some  metallic  mercury  in 
the  bottle  containing  it;  still  after  standing  some  weeks  a  basic  mercurous  nitrate 
crystallizes  out,  which  a  fresh  supply  of  nitric  acid  will  dissolve.  Sulphur 
attacks  mercury  even  in  the  barometric  vacuum,  forming  HgS  (Schrotter, 


38  MERCURY.  §58,  5J. 

J.  C.,  1873,26,  476).  H2SO4,  concentrated  at  25°  has  no  action  on  Hg  (Pitman. 
J.  Am.  Soc.,  1898,  20,  100).  With  the  hot  concentrated  acid  S02  is  evolved  and 
Hg,SO4  is  formed  if  Hg  be  in  great  excess;  HgSO4  if  the  H2SO4  be  in  excess. 
Hydrochloric  acid  gas  at  200°  is  without  action  (Berthelot,  A.  Ch.,  1856,  (3),  46, 
492);  also  the  acid  sp.  gr.,  1.20.  Bailey  and  Fowler  (J.  C.,  1888,  53,  759)  say  that 
dry  hydrochloric  acid  gas  in  presence  of  oxygen  and  mercury,  at  ordinary  tem- 
perature for  three  weeks,  forms  Hg2OCl2  without  evolution  of  hydrogen: 
^Kg-  -f  2HC1  +  (X  =  Hg2OCL,H2O  .  Hydrobromic  and  hydriodic  acids,  gases. 
both  attack  mercury,  evolve  H  ,  and  form  respectively  HgBr  and  Hgl  (Ber- 
thelot, I.  c.).  Hydrosnlphnric  acid,  dry  gas,  at  100°  does  not  attack  dry  Hg 
(Berthelot,  I.e.).  H.vdrosulphnric  acid,  in  solution,  and  alkali  sulphides  form 
HgS.  Chlorine,  bromine  and  iodine,  dry  or  moist,  attack  the  metal;  mercurous 
salts  are  formed  if  the  mercury  be  in  excess,  mercuric  palts  if  the  halogen  be  in  excess, 
b.— Oxides. — Mercurous  oxide,  Hg2O ,  is  a  black  powder  insoluble  in  water  or 
alkalis.  Hydrochloric  acid  forms  HgCl  ;  sulphuric  acid  forms  Hg2SO4  ,  changed 
by  boiling  with  excess  of  acid  to  HgSO4  ;  nitric  acid  forms  HgNO3  ,  changed  by 
excess  of  acid  to  Hg(NO3)2  •  Mercuric  oxide  is  soluble  in  acids,  insoluble  in 
alkalis,  soluble  in  20,000  to  30,000  parts  water  (Bineau,  C.  r.,  1855,  41,  509). 
It  is  red  when  produced  by  heating  in  the  dry  way  and  orange  yellow  when 
formed  by  precipitation  with  alkalis.  It  is  decomposed  by  alkali  chlorides 
forming  HgCl2*  (Mialhe,  A.  Ch.,  1842,  (3),  5,  177),  soluble  in  NH4C1  ,  from 
which  solution  NH4OH  produces  the  white  precipitate  NH4Cl,NHgH2Cl  + 
NHoHgCl  (Ditte,  C.  r.,  1891,  112,  859),  soluble  in  KI2  forming  2KI,HgI2  (Jehn, 
J.  C.,  1872,  25,  987). 

c. — Salts. — Mercury  forms  two  well  marked  classes  of  salts— mercurous. 
monovalent,  and  mercuric,  divalent — most  mercurous  compounds  are  per- 
manent in  the  air,  but  are  changed  by  powerful  oxidizing  agents  to 
mercuric  compounds.  The  latter  are  somewhat  more  stable,  but  are 
changed  by  many  reducing  agents,  first  to  merctirous  compounds  and  tben 
to  metallic  mercury  (10).  Solutions  of  mercury  salts  redden  litmus. 
Many  of  the  salts  of  mercury  are  either  insoluble  in  water,  or  require  the 
presence  of  free  acid  to  keep  them  in  solution,  being  decomposed  by  water 
at  a  certain  degree  of  dilution,  precipitating  a  basic  salt  and  leaving  an 
acid  salt  in  solution.  Mercurous  chloride,  bromide,  and  iodide  are  insolu- 
ble in  water;  the  sulphate  is  soluble  in  500  parts  cold  and  300  parts  hot 
water,  soluble  in  dilute  nitric  acid  (Wackenroder,  A.,  1842,  41,  319).  The 
acetate  has  about  the  same  solubilities  as  the  sulphate.  Mercurous  nitrate 
is  completely  soluble  in  water.  On  standing  it  gradually  changes  to 
mercuric  nitrate,  prevented  by  the  presence  of  free  mercury,  but  if  free 
mercury  be  present  a  precipitate  of  basic  mercurous  nitrate  gradually 
forms.  Mercuric  chloride  is  soluble  in  16  parts  of  cold  water  and  3  parts 

*  The  Law  of  Mass-Action  requires  that  where  the  constituents  of  a  slightly-ionized  substance 
are  present  that  substance  shall  form  at  the  expense  of  those  more  strongly  ionized.  Such  a 
slightly-ionized  body  is  HgCl2.  When  HgO  is  brought  into  contact  with  KC1  solution  Hg  and 
Cl  combine  to  form  the  non-dissociated  HgCl2,  leaving  K  and  O,  which  unite  with  water,  im- 
parting to  the  solution  a  strong  alkaline  reaction.  KBr  and  KI  act  even  more  strongly.  HgO, 
although  from  the  ready  decomposition  of  its  salts  by  water  and  from  its  easy  reducibility  a 
weak  base,  yet  will  replace  the  alkali  metals  where  a  little-dissociated  Hg  compound  results. 

An  excess  of  Hg(NO3i2  dissolves  chloride,  bromide,  and  iodide  of  Hg  and  Ag  owing  to  the 
same  cause,  the  Hg"  ions  of  the  strongly  dissociated  nitrate  decreasing  the  already  slight 
dissociatioSikof  the  mercuric  haloids  (§44).  The  failure  of  HgCla  to  give  many  of  the  pre- 
cipitation-re3||k>ns  obtainable  with  other  soluble  mercuric  salts  is  of  course  due  to  the  same 
fact— the  sligTWoncentration  of  Hg"  ions  (§  45). 


§58,  6a.  MERCURY.  39 

warm  water;  the  bromide  is  soluble  in  94  parts  water  at  9°  and  4-5  parts 
at  100°,  decomposed  by  warm  nitric  or  sulphuric  acids;  the  iodide  is 
soluble  in  about  25,000  parts  water  (Bourgoin,  A.  Cli.,  1884  (6),  3,  429), 
soluble  in  Na2S,0,  (Eder  and  Ulen,  J..C.,  1882,  42,  806),  and  in  many 
alkali  salts,  forming  double  salts.  Normal  mercuric  sulphate  is  decom- 
posed by  water  into  a  soluble  acid  sulphate  and  the  basic  sulphate,  HgS04 , 
2HgO  ,  which  is  practically  insoluble  (soluble  in  43,478  parts'  water  at 
16°,  Cameron,  Analyst,  1880,  144).  The  normal  nitrate  is  deliquescent, 
very  soluble  in  a  small  amount  of  water,  but  more  water  precipitates  the 
nearly  insoluble  basic  nitrate,  3HgO.N205 ,  changed  by  repeated  washing 
into  the  oxide,  HgO  (Millon,  A.  Ch.,  1846  (3),  18,  361).  The  basic  nitrate 
is  soluble  in  dilute  nitric  acid.  The  cyanide  is  soluble  in  eight  parts  water 
at  15°.  The  acetate  is  readily  soluble,  the  chromate  and  citrate  sparingly, 
and  the  sulphide,  iodide,  iodate,  basic  carbonate,  oxalate,  phosphate,  arse- 
nate,  arsenite,  ferrocyanide,  and  tartrate  are  insoluble  in  water. 

6.  Reactions,  a. — Fixed  alkali  hydroxides  precipitate,  from  solutions  of 
mercurous  salts,  mercurous  oxide,  Hg20 ,  black,  insoluble  in  alkalis,  readily 
transposed  by  acids;  from  solutions  of  mercuric  salts,  the  alkali,  added 
short  of  saturation,  precipitates  reddish-brown  basic  salts;  when  added  in 
excess,  the  orange-yellow  mercuric  oxide,  HgO ,  is  precipitated.  If  the 
solution  of  mercuric  salt  be  strongly  acid  no  precipitate  will  be  obtained 
owing  to  the  combination  of -the  mercuric  salt  with  the  alkali  salt  formed, 
producing  a  double  salt  in  which  the  mercury  is  present  in  the  acid  ion  un- 
affected by  the  hydroxyl  ion.  Ammonium  hydroxide  and  carbonate  pre- 
cipitate from  solutions  of.  mercurous  salts,  black  mixtures  of  mercury 
and  mercuric  ammonium  compounds.  The  same  is  true  of  the  action  of 
ammonium  h}rdroxide  on  insoluble  mercurous  salts :  2HgCl  -|-  2NH4OH 
=  Hg  +  NH,HgCl  +  2H20  -f  HH4C1  ;  6HgN03  +  6tfH4OH  =  3Hg  + 
(NH,HgNO ;)2HgO  -f  4NH4tfO!  +  5H20  ;  4Hg2S04  +  SNH4OH  =  4Hg  + 
(HgH2N)oS04.2HgO  +  3(NH4).,S04  +  6H,0  ;  or  uniting  the  salt  in  dif- 
ferent manner,  4HgCl  +  4NH4OH  ==  2Hg  -f  Hg2NCl.NH4Cl  +  2NH4C1 
-f-  4H20  .  Examination  with  a  microscope  reveals  the  presence  of  Hg°  , 
The  mercuric  ammonium  precipitate  dissolves  in  a  saturated  solution  of 
(NH4).,S04  containing  ammonium  hydroxide  and  can  thus  be  separated 
from  the  Hg  (Francois,  J.  Pliarm.,  1897  (6),  5,  388;  Turi,  Gazzctta.  1S9.S, 
23,  ii,  231;  Pesci,  Gazzetta,  1891,  21,  ii,  569;  Barfoed,  J.  pr.,  1889,  (2),  39, 
201).  With  mercuric  salts  ammonium  hydroxide  produces  "  white  precipi- 
tate/' recognizable  in  very  dilute  solutions;  that  with  cold  neutral  solu- 
tions of  mercuric  chloride  being  mercurammonium  chloride,  (NHoHg)Cl , 
also  called  nitrogen  dihydrogen  mercuric  chloride  (a)  ;  with  hot  solution 
and  excess  of  ammonium  hydroxide,  dimercurammomum  chloride, 
NHg3Cl,  also  called  nitrogen  dimercuric  chloride  (6)  is  formed,  Treat- 


40  MERCURY.  §58,  6&. 

ing  with  fixed  alkali  hydroxide  until  no  more  ammonia  is  evolved  changes 
the  former  compound  to  the  latter  (Pesci,  /.  c.).  The  precipitates  are 
easily  soluble  in  hydrochloric  acid,  slightly  soluble  in  strong  ammonium 
hydroxide,  and  more  or  less  soluble  in  ammonium  salts,  especially  am- 
monium nitrate  and  carbonate  (Johnson,  (7.  N.,  1889,  59,  23-t).  A  soluble 
combination  of  ammonium  chloride  with  mercuric  chloride,  2NH4C1. 
HgClo ,  or  ammonium  mercuric  chloride,  called  "  sal  alembroth,"  is  not 
precipitated  by  ammonium  hydroxide,  but  potassium  hydroxide  precipi- 
tates therefrom  the  white  mercurammonium  chloride,  (NH3),HgCL  (c): 

(a)  HgCl2  +  2NH4OH  =  NH2HgCl  +  NH4C1  +  2H20 
(6)  2HgCl2  +  4NH4OH  =  NHg2Cl  +  3NH4C1  +  4ELO 
(c)  2NH4Cl.HgCl2  +  2KOH  =  (NH3)2HgCL  +  2KC1  +  2H20 

A  solution  of  HgCl2  in  KI  with  an  excess  of  KOH  (Xessler's  Reagent)  is 
precipitated  by  NH4OH  (or  by  ammonium  salts),  as  NHg2I  (§207,  6k). 

Fixed  alkali  carbonates  precipitate  from  mercurous  salts  an  unstable  )ncr- 
curous  carbonate,  Hg2CO3  ,  gray,  blackening-  to  basic  carbonate  and  oxide  when 
heated.  Carbonates  of  barium,  strontium,  calcium  and  magnesium  precipitate 
mercurous  carbonate  in  the  cold.  Mercuric  salts  are  precipitated  as  red-broicn 
basic  salts,  which,  by  excess  of  the  reagent  with  heat,  are  converted  into  the 
yellow  mercuric  oxide.  The  basic  salt  formed  with  mercuric  chloride  is  an  oxy- 
chloride,  HgCl2.(HgO)2  ,  3,  or  4;  with  mercuric  nitrate,  a  basic  carbonate, 
(HgO)3HgCO3  .  Barium  carbonate  precipitates  a  basic  salt  in  the  cold,  from 
the  nitrate,  but  not  from  the  chloride. 

6. — Oxalic  acid  and  soluble  oxalates  precipitate  from  solutions  of  mercurous 
salts  mercurous  oxalatc,  Hg2C2O4  ,  white,  slightly  soluble  in  nitric  acid:  from 
solutions  of  mercuric  salts,  except  HgCl2  , '-mercuric  <t.rnlat(\  HgC204  ,  white, 
easily  soluble  in  hydrochloric  acid,  difficultly  soluble  in  nitric  acid.  A  solution 
of  HgCL  boiled  in  the  sunlight  with  (NH4),C204  gives  HgCl  and  C02  . 

Hydrocyanic  acid  and  alkali  cyanides  decompose  mercurous  salts  into  me- 
tallic mercury,  a  gray  precipitate,  and  mercuric  cyanide,  which  remains  in 
solution.  Mercuric  salts  are  not  precipitated,  since  the  cyanide  is  readily 
soluble  in  water.  Soluble  ferrocyanides  form  with  mercurous  salts  a  white  ge- 
latinous precipitate,  soon  turning  bluish  green;  with  mercuric  salts  a  white  pre- 
cipitate, soon  turning  blue.  Soluble  ferricyanides  form  with  mercurou?  salts 
a  yellowish  green  precipitate;  with  mercuric  salts  a  green  precipitate,  soluble, 
in  hydrochloric  acid.  Potassium  thiocyanate  precipitates  mercurous  thiocyanate, 
HgCNS  ,  white,  from  solutions  of  mercurous  salts  (Glaus,  J.  pr.,  1838,  15,  406)  ; 
from  solutions  of  mercuric  salts,  mercuric  thiocyanate,  Hg(CNS)2  ,  white,  soluble 
in  hot  water  (Philipp,  Z  Ch.,  1867,  553). 

c. — Nitric  acid  never  acts  as  a  precipitant  of  mercury  salts,  the  salts  being 
more  soluble  in  strong  nitric  acid  than  in  water  or  the  dilute  acid;  also  nitric 
acid  dissolves  all  insoluble  salts  of  mercury  except  HgS  ,  which  is  insoluble  in 
the  hot  acid  (sp.  yr.  1.42)  (Howe,  Am.,  1887,  8,  75).  HgCl  is  slowly  dissolved  by 
nitric  acid  on  boiling.  All  mercurous  salts  are  oxidized  to  mercuric  salts  b}' 
excess  of  nitric  acid. 

(J. — Hypophosphorous  acid  reduces  mercuric  salts  to  Hg°,  but  the  presence  of 
hydrogen  peroxide  causes  the  formation  of  HgCl  from  HgCL  and  is  of  value 
as  a  quantitative  method  for  estimation  of  mercury  (Vanino  and  Treubert,  /?., 
1897,  30,  1099). Phosphorus  acid  also  reduces  HgCl2  to  HgCl . 

Phosphoric  acid  and  alkali  phosphates  precipitate,  irom  mercurous  salts. 
mercurous  phosphate,  Hg3PO4  ,  white,  if  the  reagent  be  in  excess;  but  if  HgNO, 
be  in  excess,  Hg3PO4.HgNO3  ,  white,  with  a  yellowish  tinge.  Mercurous  phos 
phate  is  soluble  in  dilute  HNO3  ,  insoluble  in  H3P04  ,  From  mercuric  nitrate 
mercuric  phosphate,  Hg3(PO4)2,  wh'te,  h  pr3cip:tated,  soluble  in  HNO3 ,  HC1, 


§58,  6e.  M  Kit  (JURY.  41 

and  ammonium  salts,  insoluble  in  H3PO4  .  Phosphoric  acid  does  not  precipitate 
HgCl,,  and  Na.,HPO4  does  not  precipitate  the-  white  Hg-3(P04)2  from  HgCl,  , 
but  on  standing  a  portion  of  the  mercury  separates  as  a  dark  brown  pre- 
cipitate (Haack,  ,/'.  ('.,  1S91,  60,  400;  1892,  62,  5:>0). 

e. — Hydrosulphuric  acid  and  soluble  sulphides,  precipitate  from  mer- 
curous salts,  mercuric  sulphide,  HgS ,  black,  and  nie-rcury,  gray.  Mercurous 
sulphide,  Hg.JS  ,  (loos  not  exist  at  ordinary  temperatures.  According  to 
Antony  and  Sestini  (Gazzetta,  1894,  24,  i,  193),  it  is  formed  at  —  10°  by 
the  action  of  H2S  on  HgCl ,  decomposing  at  0°  into  HgS  and  Hg .  From 
mercuric  salts  there  is  formed,  first,  a  white  precipitate,  soluble  in  acids 
and  excess  of  the  mercuric  salts,  on  further  additions  of  the  reagent,  the 
precipitate  becomes  yellow-orange,  then  brown,  and  finally  black.  This 
progressive  variation  of  color  is  characteristic  of  mercury.  The  final  and 
stable  black  precipitate  is  mercuric  sulphide,  HgS  ;  the  lighter  colored 
precipitates  consist  of  unions  of  the  original  mercuric  salt  with  mercuric 
sulphide,  as  HgCl2.HgS ,  the  proportion  of  HgS  being  greater  with  the 
darker  precipitates.  When  sublimed  and  triturated,  the  black  mercuric 
sulphide  is  converted  to  the  red  (vermillion),  without  chemical  change. 
Mercuric  sulphide  is  insoluble  in  dilute  HNO,  (distinction  from  all  other 
metallic  sulphides);  insoluble  in  HC1  (Field,  J.  (7.,  1860,  12,  158);  soluble  in 
chlorine  (nitro-hydrochloric  acid);  insoluble  in  (NH4)2S  except  when  KOH 
or  NaOH  be  present  (Volhard,  A.,  1891,  255,  252);  soluble  in  K2S  (Ditte, 
C.  r.,  1884,  98,  1271),  more  readily  if  KOH  be  present  (separation  from 
Pb  ,  Ag  ,  Bi ,  and  Cu)  (Polstorff  and  Billow,  Arch.  Pharm.,  3891,  229,  292). 
A  little  HgS  (0.5-1.0  mg.)  may  dissolve  in  ammonium  polysulphide  when 
a  large  amount  of  mercury  is  present  (A.  A.  Noyes,  J.,  Am.  Cliem.  Soc., 
29,  170).  It  is  soluble  in  K2CS3  (one  part  S ,  two  parts  CS2 ,  and  23  parts 
KOH,  sp.  gr.  1.13)  (separation  from  Pb ,  Cu ,  and  Bi)  ;  reprecipitated  as 
sulphide  by  HC1  (Rosenbladt,  Z.,  1887,  26,  15). 

Mercurous  nitrate  forms  with  sodium  thiosulphate  a  grayish  black  precipi- 
tate, part  of  the  mercury  remaining  in  solution.  Mercurous  chloride  forms 
metallic  mercury  and  some  mercury  salt  in  solution  as  double  salt  (Schnauss, 
J.  (7.,  1876,  29,  342).  Mercuric  chloride  added  to  sodium  thiosulphate  forms  a 
white  precipitate,  which  blackens  on  standing1;  if  the  mercuric  chloride  be 
added  in  excess  a  bright  yellow  precipitate  is  formed,  which  blackens  when 
boiled  with  water,  nitric  acid  or  sulphuric  acid,  but  does  not  dissolve  or 
blacken  on  boiling  with  hydrochloric  acid.  Sodium  thiosulphate  added  to 
mercuric  chloride  forms  a  white  precipitate,  which  blackens  on  standing  or  on 
adding  excess  of  thiosulphate,  but  if  excess  of  thiosulphate  be  rapidty  added  to 
HgCl,  no  precipitate  is  formed;  boiling  or  long  standing  produces  the  black 
precipitate.  Mercuric  salts  are  not  completely  precipitated  by  sodium  thio- 
sulphate. The  black  precipitate  is  HgS. 

Sulphurous  acid  and  soluble  sulphites  form  from  mercurous  solutions  a 
black  precipitate  of  complex  sulphite  (Divers  and  Shimidzu,  «/.  C.,  1886,  49, 
567).  Mercuric  nitrate  with  sulphurous  acid  forms  slowly  a  flocculent  white 
precipitate  soluble  in  nitric  acid.  The  precipitate  and  solution  contain  mer- 
curosum  as  evidenced  by  HC1 .  Mercuric  nitrate  with  soluble  sulphites  forms 
a  voluminous  white  precipitate,  soluble  in  HNOS  and  containing  mercurosum. 
Mercuric  chloride  is  not  precipitated  by  sulphurous  acid  or  sulphites  in  the 
cold,  but  is  reduced,  by  boiling  with  sulphurous  acid,  to  HgCl  and  then  to  Hg° 


42  MERCURY.  §58,  6/. 

Sulphuric  acid  and  soluble  sulphates  precipitate  from  mercurous  solu- 
tions not  too  dilute,,  mercurous  sulphate,  Hg2S04 ,  white,  decomposed  by 
boiling  water,  sparingly  soluble  in  cold  water  (5c),  soluble  in  nitric  acid 
and  blackened  by  alkalis.  Mercuric  salts  are  not  precipitated  by  sulphuric 
acid  or  sulphates.  For  action  of  H2S04  on  HgCl2  see  next  paragraph  and 
(§269,  8,  footnote). 

/. — Hydrochloric  acid  and  soluble  chlorides  precipitate  from  solutions  of 
mercurous  salts,  mercurous  chloride,  HgCl ,  "  Calomel,"  white,  insoluble  in 
water,  slowly  soluble  in  hot  concentrated  HC1 .  Boiling  nitric  acid  slowly 
dissolves  it,  forming  Hg(N"03)2  and  HgCL  ;  dissolved  by  chlorine  or  nitro- 
hydrochloric  acid  to  HgCl2  ;  soluble  in  Hg(NO.,)2  (57;  footnote)  (Dresehspi, 
J.  C.,  1882,  42,  18).  This  precipitation  of  mercurous  salts  by  hydro- 
chloric acid  is  a  sharp  separation  from  mercuric  salts  and  places  mer- 
curous mercury  in  the  FIRST  (SILVER)  GROUP  OF  METALS.  Mercuric  salts 
are  not  precipitated  by  hydrochloric  acid  or  soluble  chlorides,  unless  the 
mercuric  solution  is  more  concentrated  than  possible  for  a  mercuric 
chloride  solution  under  the  same  conditions,  i.  e.,  a  strong  solution  of 
Hg(N03)2  gives  a  precipitate  of  HgCl2  on  addition  of  HC1 ,  soluble  on 
addition  of  water.  Mercuric  chloride  is  not  decomposed  by  sulphuric 
acid.  A  compound  HgCl2.H2S04  is  formed  which  sublimes  undecom- 
posed.  The  same  compound  is  formed  when  HgS04  is  treated  with  HC1 
and  distilled  (Ditte,  A.  Ch.,  1879,  (5),  17,  120). 

Hydrobromic  acid  and  soluble  bromides  precipitate,  from  solutions  of 
mercurous  salts,  mercurous  bromide,  HgBr ,  yellowish  white,  insoluble  in 
water,  alcohol,  and  dilute  nitric  acid;  from  concentrated  solutions  of 
mercuric  salts,  mercuric  bromide,  HgBr2 ,  white,  decomposed  by  concen- 
trated nitric  acid.  Mercuric  bromide  is  soluble  in  excess  of  mercuric  salts 
(5&  footnote),  or  in  excess  of  the  precipitant;  hence,  unless  added  in 
suitable  proportions,  no  precipitate  will  be  produced.  Sulphuric  acid  does 
not  transpose  HgBr2  but  forms  compounds  exactly  analogous  to  those 
with  HgCl2 .  Excess  of  concentrated  H2S04  gives  some  Br  with  HgBr2 . 

Hydriodic  acid  and  soluble  iodides  precipitate  from  solutions  of  mer- 
curous salts,  mercurous  iodide,  Hgl,  greenish  yellow — "the  green  iodide 
of  mercury" — nearly  insoluble  in  water,  insoluble  in  alcohol  (distinction 
from  mercuric  iodide),  soluble  in  mercurous  and  mercuric  nitrates;  decom- 
posed by  soluble  iodides  with  formation  of  Hg  and  HgI2 ,  the  latter  being 
dissolved  as  a  double  salt  with  the  soluble  iodide:  2HgI  4-  2KI  —  Hg  + 
HgI2.2KI .  Mercurous  chloride  is  transposed  by  HI  or  KI  to  form  Hgl , 
excess  of  the  reagent  reacts  according  to  the  above  equation  (D.,  2,  2,  867). 
Ammonium  hydroxide  in  the  cold  decomposes  Hgl  into  Hg  and  HgI2 
(Francois,  J.  Pharm.,  1897,  (6),  5,  388). 

Mercuric  salts  are  precipitated  as  mercuric  iodide,  HgI2 ,  first  reddish- 


§1)8,  7.  MERCURY.  43 

yellow  then  red,  soluble  in  24,814  parts  of  water  at  17.5°  (Bourgoin,  A.  CU., 
1884,  (6),  3,  439),  soluble  in  concentrated  nitric  and  hydrochloric  acids; 
quickly  soluble  in  solutions  of  the  iodides  of  all  the  more  positive  metals, 
i.  e.  in  excess  of  its  precipitant,  by  formation  of  soluble  double  iodides;  as 
(KI).,HgI.,  variable  to  KIHgI2 .  A  hot  concentrated  solution  of  potas- 
sium iodide  dissolves  3HgI2  for  every  2KI .  The  first  crystals  from  this 
solution  are  KIHgl., .  These  are  decomposed  by  pure  water,  and  require 
a  little  alkali  iodide  for  perfect  solution,  but  they  are  soluble  in  alcohol 
and  ether.  A  solution  of  dipotassium  mercuric  tetraiodide,  K2HgI4  = 
(KI),HgI,  (sometimes  designated  the  iodo-hydrargyrate  of  potassium),  is 
precipitated  by  ammonium  hydroxide  as  mercurammonium  iodide,  NHg2I 
( Xcssler's  test),  and  by  the  alkaloids  (Mayer's  reagent). 

Potassium  bromate  precipitates,  from  solutions  of  mercurous  nitrate,  mer- 
curous  bromate,  HgBr03  ,  white,  soluble  in  excess  of  mercurous  nitrate  and 
in  nitric  acid;  from  solutions  of  mercuric  nitrate,  mercuric  bromate,  Hg(Br03)2, 
white,  soluble  in  nitric  acid,  hydrochloric  acid,  and  in  excess  of  mercuric  nitrate, 
in  650  parts  of  cold  and  64  parts  of  hot  water  (Rammelsberg,  Pogg.,  1842, 
55,  79).  No  precipitate  is  formed  when  potassium  bromate  is  added  to  mercuric 
chloride  (5&,  footnote).  lodic  acid  and  soluble  iodates  precipitate  solutions 
of  mercurous  salts  as  mwcurous  iodate,  HgI03  ,  white  with  yellowish  tint,  solu- 
ble with  difficulty  in  dilute  nitric  acid,  readily  soluble  in  HC1  by  oxidation  to 
mercuric  salt.  Mercuric  nitrate  is  precipitated  as  mercuric  iodate,  Hg(I03)2  , 
white,  soluble  in  HC1 ,  insoluble  in  HN03  and  H2S04  ,  soluble  in  NH4C1  ,  trans- 
posed and  then  dissolved  by  KI .  Mercuric  chloride  is  not  precipitated  by 
KIO3  (56,  footnote)  (Cameron,  C.  N.,  1876,  33,  253). 

g. — Arsenous  acid  or  arsenites  form  a  white  precipitate  with  mercurous 
nitrate,  soluble  in  HN03  (Simon,  Pogg.,  1837,  40,  442).  Mercuric  nitrate  is 
precipitated  by  a  solution  of  arsenous  acid;  the  precipitate  is  soluble  in  HNO3 
(/>.,  2,  2,  920).  Arsenic  acid  or  Na,HAsO4  precipitates  from  mercurous  nitrate 
:{Hg3As04.HgNO3.H20  ,  lig-ht  yellow  if  the  HgNO3  be  in  excess  (D.,  2,  2,  921); 
dark  red  Hg3As04  if  the  arsenate  be  in  excess.  Hg3AsO4  is  changed  by  cold 
HC1  to  HgCl  and  H3As04  ,  by  boiling-  with  HC1  to  Hgo  ,  HgCl,  ,  and  H3AsO4 ; 
and  is  soluble  unchanged  in  cold  HNO3  ,  insoluble  in  water  and  acetic  acid 
(Simon,  Pogg.,  1837,  41,  424).  Arsenic  acid  and  soluble  arsenates  precipitate 
from  mercuric  nitrate,  Hg3(AsO4)2  ,  white,  soluble  in  HNO3  and  HC1 ,  slightly 
soluble  in  water.  Arsenic  acid  and  potassium  arsenate  do  not  precipitate 
mercuric  chloride  from  its  solutions. 

Stannous  chloride  precipitates  solutions  of  mercuric  salts  (by  reduction), 
as  mercurous  chloride,  white;  or  if  the  stannous  chloride  be  in  excess, 
as  metallic  mercury,  gray  to  black  (a  valuable  final  test  for  mercuric  salts)  (10). 

ft. — Soluble  chromates  precipitate  from  mercurous  solutions  mercurous 
tfu'oniate,  Hg,CrO4  ,  brick-red,  insoluble  in  water,  readily  transposed  by  HC1  to 
HgCl  and  H2Cr04  ,  soluble  with  difficulty  in  HNO3  without  oxidation  (Richter, 
B.,  1882,  15,  1489).  Mercuric  nitrate  is  precipitated  by  soluble  chromates  as  a 
light  yellow  precipitate,  rapidly  turning  dark  brown,  easily  soluble  in  dilute 
acids  and  in  HgCl,.  Mercuric  chloride  forms  a  precipitate  with  normal  chro- 
mates, but  not  with  KoCr.,O7  . 

7.  Ignition. — Mercury  from  all  its  compounds  is  volatilized  by  heat  a* 
the  undecomposed  salt  or  as  the  free  metal.  Mercurous  chloride  (Debray, 


44  MERCURY.  §58,  8. 

J.  C.,  1877,  31,  47)  and  bromide  and  mercuric  chloride  and  iodide  sublime 
(in  glass  tubes)  undecomposed — the  sublimate1  condensing  (in  the  cold  part 
of  the  tube)  without  change.  Most  other  compounds  of  mercury  are 
decomposed  by  vaporization,  and  give  a  sublimate  of  metallic  mercury 
(mixed  with  sulphur,  if  from  the  sulphide,  etc.).  All  compounds  of  mer- 
cury, dry  and  intimately  mixed  with  dry  sodium  carbonate,  and  heated  in 
a  glass  tube  closed  at  one  end,  give  a  sublimate  of  metallic  mercury  as  a 
gray  mirror  coat  on  the  inner  surface  of  the  cold  part  of  the  tube.  Under 
the  magnifier,  the  coating  is  seen  to  consist  of  globules,  and  by  gently 
rubbing  with  a  glass  rod  or  a  wire,  globules  visible  to  the  unaided  eye  are 
obtained. 

8.  Detection. — Mercury  in  the  mercurous  condition  belongs  to  the  FIRST 
GROUP  (silver  group),  and  is  completely  precipitated  by  HC1 .     It  is  iden- 
tified by  the  action  of  ammonium  hydroxide,  changing  the  white  precipi- 
tate of  mercurous  chloride  to  the  black  precipitate  of  metallic  mercury 
and  nitrogen  dihydrogen  mercuric  chloride  (a  delicate  and  characteristic 
test  for  Hg').     Mercury  in  the  mercuric  condition  belongs  to  the  SECOND 
GROUP  (tin  and  copper  group),  and  is  separated  from  all  other  metals  of 
that  group  by  the  non-solubility  of  the  sulphide  in  (NH4)2SX  and  in  dilute 
HN03  .    The  sulphide  is  dissolved  in  nitrohydrochloric  acid,  and  the  pres- 
ence of  mercury  confirmed  by  the  precipitation  of  Hg°  on  a  copper  wire,  or 
by  the  reduction  to  HgCl  or  Hg°  by  SnCl2 . 

9.  Estimation. — (a)  As  metallic  mercury.     The  mercury  is  reduced  by  means 
of  CaO  in  a  combustion-tube  at  a  red  heat  in  a  current  of  CO2  .     The  sublimed 
mercury  is  condensed  in  a  flask  of  water,  and,  after  decanting-  the  water,  dried 
in  a  bell-jar  over  sulphuric  acid  without  application  of  heat.     The  mercury  may 
also  be  reduced  from  its  solution  by  SnCL    (or  H3P03  at  100°)    and  dried  as 
above,     (ft)  As  mercurous  chloride.     It  is  first  reduced  to  Hg'  by  H3PO3   (Uslar, 
Z.,  1895,  34,  391),  which  must  not  be  heated  above  60°,  otherwise  metallic  mer- 
cury will  be  formed;  and  after  precipitation  by  HC1  and  drying-  on  a  weighed 
filter  at  100°,  it  is  weighed  as  HgCl  .     Or  enough  HC1  is  added  to  combine  with 
the  mercury,  then  the  Hg"  is  reduced  to  Hg'  by  FeSO4  in  presence  of  NaOH : 
2HgO  +  2FeO  +  3H2O  =  Hg,,0  +  2Fe(OH)s.    H2SO4  is  added,  which  causes  the 
formation  of  HgCl  ,  which  is  dried  on  a  weighed  filter  at  100°.     (c)  As  HgS  . 
It  is  precipitated  by  H2S,  and  weighed  in  same  manner  as  the  chloride.     Any 
free  sulphur  mixed  with  the  precipitate  should  be  removed  by  CS2  .     (d)    As 
HgO  ,  by  heating  the  nitrate  in  a  bulb-tube  in  a  current  of  dry  air  not  hot 
enough   to  decompose    the   HgO.     (e)    Volumetrically,   by    Na2S203;    from    the 
nitrate  the  precipitate  is  yelloiL-.  from  the  chloride  it  is  white: 

3Hg(N03)2  +  2Na2S203  +  2H2O  =  Hg3S2(NO3)2  +  2Na,SO4  +  4HNO3 
SHgCl,  -f  2Na2S203  +  2H20  —  Hg3S2CL  -f  2N;a2S04  -f  4HC1  . 

(f)  Volumetrically,  HgCL  is  reduced  to  Hg20  by  FeSO4  in  presence  of  KOH  , 
and  after  acidulating-  with  H2SO4  the  excess  of  FeS04  is  determined  by  K2Cr2O- 
or  KMnO4  (Jliptner,  C.  C.,  1882,  727).  (#)  By  iodine.  It  is  converted  into  HgCl 
and  then  dissolved  in  a  graduated  solution  of  I  dissolved  in  KI :  2HgCl  +  OKI  + 
Ia  =  2K2HgI4  +  2KC1  .  The  excess  of  iodine  is  determined  by  Na,,S,O3  .  (7i) 
The  measured  solution  of  HgCL,  is  added  to  a  graduated  solution  of  KI: 
4KI  -f-  HgCL  =  XoHgl.  +  2KC1  .  The  instant-  the  amount  of  HgCL  shown 
in  the  equation  is  exceeded  a  red  precipitate  of  HgI2  appears,  (t)  Volumetric, 


§59, 2.  SILVER.  45 

by  adding  a  few  drops  of  ammonium  hydroxide  to  HgCl,  and  then  titrating 
with  standard  KCN  ,  the  ammonium  hydroxide  precipitate  disappears  when  the 
mercury  becomes  Hg(CN)2  (Haimay,  J.  C.,  187H,  26,  570;  Tuson,  J.  C1.,  1877,  32, 
679).  (;)  Electrolytwally,  by  obtaining  the  mercury  as  HgNO3  ,  Hg(NO3)2  , 
or  Hg2SO4  and  precipitating  as  Hg°  on  platinum  by  the  electric  current. 
Mercuric  chloride  cannot  be  used,  as  it  is  partly  reduced  to  HgCl ,  and  that 
is  not  readily  reduced  to  Hg°  by  the  electric  current  (Hannay,  I.e.). 

10.  Oxidation. — Free  mercury  (Hg°)  precipitates  Ag,  An ,  and  Pt  from 
their  solutions,  and  reduces  mercuric  salts  to  mercurous  salts  (Hada,  J.  C., 
1896,  69,  1667).  Potassium  permanganate  in  the  cold  oxidizes  the  metal 
to  Hg20  ,  when  hot  to  HgO  (Kirchmann,  J.  C.,  1873,  26,  476).  Mercury 
and  mercurous  salts  are  oxidized  to  mercuric  salts  by  Br ,  Cl ,  I ,  HN03 , 
H2S04  (concentrated  and  hot),  and  HC10;5 . 

Reducing  agents,  as  Pb ,  Sn ,  Sn",  Bi ,  Cu  ,  Cu',  Cd ,  Al ,  Fe ,  Co ,  Zn  , 
Th1,  Mg,  H3P02 ,  H3P03  and  H2S03 ,  precipitate,  from  the  solutions  of 
mercuric  and  mercurous  nitrates,  dark-gray  Hg°  ;  from  solution  of  mer- 
curic chloride,  or  in  presence  of  chlorides,  first  the  white.,  HgCl ,  then  gray 
Hg°.  Strong  acidulation  with  nitric  acid  interferes  with  the  reduction, 
and  heating  promotes  it. 

The  reducing  agent  most  frequently  employed  is  stannous  chloride: 

2HgCL  +  SnCl2  =  2HgCl  +  SnCl, 

2HgCl  -f  SnCL  =  2Hg  +  SnCl4 
or  HgCL  +  SnCl2  =  Hg  +  SnCl4 

also  2Hg(N03)2  +  SnCl2  —  2HgCl  +  Sn(N03)4 

A  clean  strip  of  copper,  placed  in  a  slightly  acid  solution  of  a  salt  of  mer- 
cury, becomes  coated  with  metallic  mercury,  and  when  gently  rubbed 
with  cloth  or  paper  presents  the  tin-white  lustre  of  the  metal;  the  coating 
being  driven  off  by  heat;  2HgN03  +  Cu  =  2Hg  +  Cu(N03)2 .  Formic  acid 
reduces  mercuric  to  mercurous  chloride,  and  in  the  cold  does  not  affect 
further  reduction.  Dry  mercuric  chloride,  moistened  with  alcohol,  is 
reduced  by  metallic  iron,  a  bright  strip  of  which  is  corroded  soon  after 
immersion  into  the  powder  tested  (a  delicate  distinction  from  mercurous 
chloride). 

§59.  Silver  (Argentum)  Ag  =  107.88.      Monovalent. 

1.  Properties. — Specific  (/rarity  10.512  heated  in  vacuo  (Dumas,  C.  N.,  1878,  37, 
82).     Melting  point,  960.7°  (Heycock  and  Neville,  /.  C.,  1895,  67,  1024).     Does  not 
appreciably  vaporize  at  1567°   (V.  and  C.  Meyer,  B.,  1879,   12,  1428).     It  is  the 
whitest  of  metals,  harder  than  gold  and  softer  than  copper.     Silver  is  hardened 
by  copper;  United  States  silver  coin  contains  90  per  cent  silver  and  10  per  cent 
copper.     In  malleability  and  ductility  it  is  inferior  only  to  gold;  and  as  a  con- 
ductor of  heat  and  electricity  it  exceeds  all  other  metals. 

2.  Occurrence. — Found  in  a  free  state  in  United  States,  Mexico,  Peru,  Siberia, 
etc.;    alone,  and  with  gold  as  a  component  of  other  minerals,  e.  g.,  galena,  pyrite, 
phalcopyrite,    and   many   other   ores.     The    most   important   silver   minerals  are 

iRejd,  c,  N,,  1865,  12,  242;  Neumann,  J.  £..,  J875,  38,  J32, 


CALIFORNIA  COLLEGE 


46  SILVEE.  §59,  3. 

argentite,  Ag2S,    stephanite,  Ag5SbS4 ,    pyrargyrite,  Ag3SbS3,  proustite,  Ag3AsS3 , 
cerargyrite,  AgCl  . 

3.  Preparation. — (a)  Argentiferous  ores  are  smelted  with  lead  ores,  coke  and 
limestone  in  a  blast  furnace;    silver  (and  gold)  alloys  with  the  reduced  lead,  and 
is    subsequently    separated    from    it    by    Parkes'    or    Betts'    process.      (6)   It  is 
amalgamated  with  mercury  and   the   mercury  separated  by  distillation,     (c)    It 
is  brought  into  solution  and  the  metal  precipitated  by  copper,      (d)  It   is  very 
easily  reduced  from  the  oxide  or  carbonate  by  heat  alone,  and  from  all  its  [com- 
pounds by  ignition  with  hydrogen,  carbon,  carbon  monoxide  and  organic  compounds. 

4.  Oxides. — Silrer  oxide,  Ag2O  ,  argentic  oxide,  is  formed  by  the  action  of 
alkali  hydroxides  on  silver  salts  or  by  heating1  the  carbonate  to  200°.     It  is  a 
brown  powder,  a  strong  oxidizing  agent,  decomposed  at  300°  into  metallic  silver 
and  oxygen.     Concerning  the  existence  of  argentous  oxide,  Ag4O  ,  and  silver 
peroxide,  Ag*.>O2  ,  and  their  properties,  see  Muthmann  (J5.,  1887,  20,  983) ;  Pford- 
ten  (B.,  1887,  20,  1458)  and  Bailey  (C.  N.,  1887,  55,  263). 

5.  Solubilities. — a. — Metal. — The  fixed  alkali-s  do   not   act  upon    silver,  hence 
silver  crucibles  are  used  instead  of  platinum  for  fusion  with  caustic  alkalis. 
Ammonium  hydroxide  dissolves  finely  divided  silver,  no  action  if  air  be  excluded. 
Acetic  acid  is  without  action  (Lea,  Am.  £.,  1892,  144,  444).     Nitric  acid  is  the 
ordinary  solvent  for  silver,  the  50  per  cent  acid  being  most  effective,  while  the 
dilute  acid  free  from  nitrous  acid  has  little  or  no  action    (Lea,   I.   c.);    silver 
nitrate  is  formed,   nitrogen  peroxide  being  the  chief  product  of  the  reduction 
of  the  nitric  acid  (Higley  and  Davis,  Am.,  1897,  18,  587).     Silver  is  not  oxidized 
by  water  or  air  at  any  temperature;    it  is  attacked  by  phosphorus  or  by  sub- 
stances easily  liberating  phosphorus;    it  is  tarnished  in  contact  with  hydrosul- 
phuric  acid,  soluble  sulphides,  and  many  organic  compounds  containing  sulphur; 
except  that  pure  dry  hydrosulphuric  acid  is  without  action  upon  pure  dry  silver 
(Cabell,    C.   N.,    1884,   50,   208).     Dilute  sulphuric   acid  slowly   dissolves   finely 
divided  silver  (Lea,  1.  c.),  a  sulphate  being  formed  while,  with  the  hot  concen- 
trated acid,  sulphur  dioxide  is  evolved.     Hydrochloric  acid,  sp.  gr.,  1.20,  is  without 
action  upon  pure  silver,  but  the  metal  is  readily  attacked  by  chlorine,  bromine 
or  iodine,     b. — Oxide. — Silver  oxide,  Ag2O  ,  soluble  in  3000  parts  of  water,  com- 
bines with  nearly  all  acids,  except  CO2  ,  forming  the  corresponding  salts.     The 
hydroxide  is  not  known. 

c. — Sails. — Silver  forms  a  greater  number  of  insoluble  salts  than  any 
other  known  metal,  though  in  this  respect  mercury  and  lead  are  quite 
similar.  The  nitrate  is  very  soluble  in  water,  100  parts  water  dissolv- 
ing 227.3  parts  at  19.5°,  soluble  in  glycerol,  and  sparingly  soluble  in 
alcohol  and  ether.  The  chlorate  dissolves  in  about  ten  parts  cold  water; 
the  acetate  in  100  parts;  the  sulphate  in  about  200  parts  cold  water  and 
88  parts  at  10G°,  and  is  more  soluble  in  nitric  or  sulphuric  acid  than  in 
water;  the  borate,  thiosulphate,  and  citrate  are  sparingly  soluble  in  water. 
The  oxalate,  tartrate,  carbonate,  cyanide,  ferrocyanide,  ferricyanide,  phos- 
phate, sulphide,  sulphite,  chloride,  bromide,  iodide,  iodate,  arsenite,  arse- 
nate,  and  chromate  are  insoluble  in  water. 

The  chloride  is  soluble  in  244  parts  HC1 ,  but  its  solubility  is  very  much 
lessened  by  the  presence  of  mercurous  chloride  (Ruyssen  and  Varenne,  Bl.3 
1881,  36,  5).  If  a  solution  of  silver  nitrate  be  dropped  into  concentrated 
hydrochloric  acid  no  precipitate  appears  until  one  half  per  cent  of  the 
HC1  becomes  AgCl  (Pierre,  J.  C.,  1872,  25,  123).  Concentrated  nitric  acid 
upon  long  continued  boiling  scarcely  attacks  AgCl  (Thorpe,  J.  C.,  1872,  25, 
453);  sulphuric  acid,  sp.  gr.  1.84,  completely  transposes  even  the  fused 


;<59, 66.  SILVER.  47 

chloride  on  long  boiling  (Sauer,  J.  C.}  1874,  27,  335).  Silver  chloride  is 
also  soluble  in  ammonium  hydroxide  and  carbonate;  in  sodium  chloride 
forming  a  double  salt;  in  a  concentrated  solution  of  mercuric  nitrate 
(§68,  1;  §58,  56  footnote);  and  in  many  other  metallic  chlorides  and 
alkali  salts  to  a  greater  or  less  extent.  All  the  salts  of  silver  which  are 
insoluble  in  water  are  soluble  in  ammonium  hydroxide,  except  the  sulphide 
and  iodide;  in  ammonium  carbonate,  except  the  bromide,  iodide,  and 
sulphide,  the  bromide  very  slightly  soluble;  in  cold  dilute  nitric  acid, 
except  the  chloride,  bromide,  bromate,  iodide,  iodate,  cyanide,  and  thio- 
c van ate;  in  a  solution  of  potassium  cyanide  (and  by  many  other  cyanides) 
except  the  sulphide;  and  in  alkali  thiosulphates  almost  without  exception. 

6.  Reactions,  a. — The  fixed  alkali  hydroxides  precipitate  from  solu- 
tions of  silver  salts  (in  absence  of  citrates),  silver  oxide,  Ag20 ,  grayish 
brown,  insoluble  in  excess  of  the  reagents;  soluble  in  acids,  alkali  cyanides, 
and  thiosulphates;  somewhat  soluble  in  ammonium  salts.  Most  silver 
salts  are  transposed  011  boiling  with  the  fixed  alkalis,  except  the  iodide, 
which  is  not  thus  transposed  (Vogel,  J.  C.,  1871,  24,  313). 

Ammonium  hydroxide,  in  neutral  solutions  of  silver  salts,  forms  the 
same  precipitate,  Ag,0 ,  very  easily  dissolving  in  excess,  by  formation  of 
silver  ammonium  hydroxide,  NH3AgOH  :  AgN03  +  2NH4OH  =  NH8AgOH 
+  NH4N03  +  H20  (Prescott,  J.  Am.  Soc.,  1880,  2,  32).  In  solutions  con- 
taining much  free  acid,  all  precipitation  is  prevented  by  the  ammonium  salt 
formed  with  the  formation  of  silver  ammonium  nitrate,  NH3AgN03  or  in 
the  presence  of  excess  of  ammonia  as  (NH3).,AgN03 . 

Alkali  carbonates  precipitate  silver  carbonate,  Ag2C03 ,  white  or  yellow- 
ish white,  very  slightly  soluble  in  water  and  in  the  fixed  alkali  carbonates, 
readily  soluble  in  ammonium  hydroxide  -and  carbonate,  transposed  by 
inorganic  acids  forming  the  corresponding  salts.  Carbon  dioxide  does 
not  transpose  silver  salts. 

&. — Oxalic  acid  and  soluble  oxalates  precipitate  silver  oxalatc,  Ag,,C2O4  ,  white, 
slightly  soluble  in  water,  soluble  with  difficulty  in  dilute  nitric  or  sulphuric 
acids,  readily  soluble  in  ammonium  hydroxide.  When  heated  it  decomposes 
with  detonation,  forming-  metallic  silver. 

Potassium  cyanide  precipitates  from  neutral  or  slightly  acid  solutions 
jilrrr  cyanide.,  AgCN ,  white,  quickly  soluble  in  excess  of  the  reagent  as 
silver  potassium  cyanide,  AgCN.KCN .  Hydrocyanic  acid  precipitates 
solutions  of  silver  salts  but  the  precipitate  does  not  dissolve  in  excess  of 
the  reagent.  Silver  cyanide  is  transposed  by  HL,S04  or  HC1  and  is  soluble 
in  ammonium  hydroxide  and  carbonate  (Schneider,  J.  pr.,  18(58,  104,  83). 
The  ready  solubility  of  nearly  all  silver  compounds  in  potassium  cyanide 
(5r)  affords  a  means  of  separating  silver  from  many  minerals. 

Potassium  ferrocyanide  precipitates  silver  ferrooyanide,  Ag4Fe(CN)e,  yellow- 
ish white,  soluble  with  difficulty  in  ammonium  hydroxide  and  carbonate; 


48  SILVER.  §59, 6c. 

metallic  silver  separates  on  boiling"  and  a  ferricyanide  is  formed.  The  ferro- 
cyanide  is  not  decomposed  by  hydrochloric  acid,  but  it  is  changed  to  the 
ferricyanide  by  nitric  acid.  Exposure  to  the  air  gives  it  a  blue  tinge.  Potas- 
sium ferricyanide  precipitates  xiln-r  fcrrici/anide,  Ag3Fe(CN)6  ,  reddish  yellow, 
readily  soluble  in  ammonium  hydroxide  and  carbonate.  Potassium  thlocyanate 
gives  silver  thiocyanate,  AgCNS  ,  white,  soluble  in  ammonium  hydroxide  and 
carbonate,  insoluble  in  dilute  acids.  Concentrated  sulphuric  acid  with  the  aid 
of  heat  dissolves  silver  thiocyanate  when  some  free  silver  nitrate  is  present.  This 
may  be  used  as  a  separation  from  silver  chloride,  which  is  transposed  by  hot 
concentrated  sulphuric  acid  only  on  long-continued  boiling  (5c).  To  effect  this 
separation  a  little  silver  nitrate  should  be  added  to  the  silver  precipitates  and 
then  concentrated  sulphuric  acid  and  heat.  To  avoid  danger  of  decomposition 
of  the  chloride  the  mixture  should  not  be  heated  above  200°.  The  pure  silver 
thiocyanate  (silver  nitrate  being  absent)  is  decomposed  by  hot  concentrated 
sulphuric  acid  with  formation  of  a  black  precipitate  containing-  silver. 

c. — Silver  nitrate  is  soluble  in  500  parts  of  concentrated  nitric  acid  (Schultz, 
Z.  Ch.,  1869,  531),  and  is  precipitated  from  its  concentrated  water  solutions  by 
the  addition  of  concentrated  nitric  acid.  <1. — Disodium  phosphate  precipitates 
silver  phosphate,  Ag-3P04  ,  yellow,  soluble  in  dilute  nitric  acid,  in  phosphoric 
acid,  and  in  ammonium  hydroxide  and  carbonate;  but  little  soluble  in  dilute 
acetic  acid.  Sodium  pyrophosphate  precipitates  silver  /)//ro/>/f06'y)7/«£e,  white,  same 
solubilities  as  the  orthophosphate. 

e. — Hydrosulphuric  acid  and  soluble  sulphides  precipitate  from  neutral, 
acid  or  alkaline  solutions  silver  sulphide.,  Ag.,8 ,  black,  soluble  in  moderately 
strong  nitric  acid  (distinction  from  mercur}*),  slightly  soluble  in  potassium 
cyanide  (distinction  from  copper),  insoluble  in  alkali  sulphides  (distinction 
from  arsenic,  antimony,  and  tin).  Certain  insoluble  sulphides  form  silver 
sulphide  from  solutions  of  silver  nitrate,*  e.  g.,  cupric  sulphide  gives  silver 
sulphide,  cuprous  sulphide  gives  silver  sulphide  and  metallic  silver,  in  both 
cases  cupric  nitrate  resulting  (Schneider,  J.  C.t  1875,  28,  133  and  612). 

Thiosulphates  precipitate  silver  thiosulphate,  Ag?S2O3  ,  white,  unstable,  readily 
soluble  in  excess  of  the  precipitant,  by  formation  of  double  thiosulpahtes; 
with  excess  of  sodium  thiosulphate  Na4Ag2(S2O3)3  is  formed  (Cohen,  J.  C.,  1896, 
70,  ii,  167).  Silver  thiosulphate  turns  black  on  standing  or  heating;  Ag;SLO  + 
H2O  =  Ag2S  +  H2SO4 .  Sulphurous  acid  and  soluble  sulphites  precipitate 
silver  sulphite,  Ag2SO3  ,  white,  readily  soluble  in  excess  of  alkali  sulphite  or  in 
dilute  nitric  acid;  on  boiling  precipitated  as  metallic  silver  with  formation  of 
sulphuric  acid.  Sulphuric  acid  and  soluble  sulpht'ies  precipitate  silver  sul- 
phate, Ag.jSO.1  ,  white,  from  concentrated  solutions  of  the  nitrate  or  chlorate; 
sparingly  soluble  in  water,  quite  soluble  in  concentrated  sulphuric  acid. 

/. — Hydrochloric  acid  and  soluble  chlorides  precipitate  silver  chloride, 
AgCl ,  white,  curdy ;  separated  on  shaking  the  solution ;  turning  violet  to 
brown  on  exposure  to  the  light;  fusible  without  decomposition;  very 
easily  soluble  in  ammonium  hydroxide  as  ammonia  silver  chloride, 
(NH3)3(AgCl)2  (Jarry,  C.  r.,  1897,  124,  288),  according  to  the  following 

equation :  f 

2AgCl  +  3NH4OH  -  3NH3 .2AgCl  +  3H2O. 

On  acidifying  the  solution  with  nitric  acid  the  silver  chloride  is  repre- 
cipitated  as  follows : 

3NH3 .2AgCl  +  3HNO3  =  2AgCl  +  3NH4NO3. 

If  mercurous  chloride  be  present  with  silver  chloride  the  solubility  in  ammo- 
nium hydroxide  is  greatly  lessened,  in  fact  a  great  excess  of  mercurous 

*  Ag2$  is  one  of  the  least  soluble  of  the  sulphides.     See  §  57,  6e,  footnote. 


§59,  7.  SILVE1L  49 

chloride  may  entirely  prevent  the  solution  of  silver  chloride  in  ammonium 
hydroxide  by  forming  metallic  silver.  AgCl  +  :$HgCl  +  4NH4OH  = 
Ag  +  2Hg>  +  2NH2HgCl  +  2NH4C1  +  4H20.  Silver  chloride  is  quite 
soluble  in  a  solution  of  mercuric  nitrate,  which,  if  present  in  large  excess, 
may  entirely  prevent  the  precipitation  of  the  silver  chloride  by  hydrochloric 
acid.  The  precipitation  by  hydrochloric  acid  (in  absence  of  a  great  excess 
of  Hg(NO;i)2)  is  the  most  delicate  of  the  ordinary  tests  for  silver,  being 
recognized  in  250,000  parts  of  water.  As  mercuric  salts  are  not  at  all  pre- 
cipitated by  HC1  and  lead  salts  only  imperfectly,  silver  is  the  only  metal 
which  belongs  exclusively  to  the  FIRST  OR  SILVER  GROUP  OF  BASES  (§16). 

Hydrobromic  acid  and  soluble  bromides  precipitate  silver  bromide,  AgBr , 
white,  with  a  slight  yellowish  tint;  but  slightly  soluble  in  excess  of  alkali 
bromides,  and  much  less  easily  soluble  in  ammonium  hydroxide  than  silver 
chloride.  If  silver  nitrate  be  added  to  a  bromide  containing  an  excess  of  am- 
monium hydroxide,  the  precipitate  which  first  forms  readily  dissolves  on  shak- 
ing", no  solution  is  obtained  with  the  iodide. 

Hydriodic  acid  and  soluble  iodides  precipitate  silver  iodide,  Ag-I  ,  pale  yellow, 
soluble  in  excess  of  the  concentrated  reagents  by  formation  of  double  iodides, 
as  KIAgT  ,  which  are  decomposed  by  dilution  with  much  water.  The  precipi- 
tate dissolves  in  26,000  parts  of  ten  per  cent  ammonium  hydroxide;  not  at  all  in 
a  five  per  cent  solution  (Longi,  Gazzetta,  1883,  13,  87).  It  is  insoluble  in  dilute 
acids,  but  is  decomposed  by  hot  concentrated  nitric  or  sulphuric  acids. 

Silver  bro-mate  formed  by  adding-  potassium  bromate  to  silver  nitrate  is  soluble 
in  about  600  parts  water  and  in  320.4  parts  nitric  acid  (sp.  gr.t  1.21)  at  25°,  and 
readily  soluble  in  ammonium  hydroxide.  Silver  iodate  formed  in  manner  simi- 
lar to  the  bromate  is  soluble  in  about  28,000  parts  water  and  in  1044.3  parts 
nitric  acid  (sp.  gr.,  1.21)  at  25°,  and  readily  soluble  in  ammonium  hydroxide 
(Longi,  I.e.). 

0.— Soluble  arsenites  precipitate  silver  arsenite,  Ag3As03  ,  yellow,  very  readily, 
soluble  in  dilute  acids  and  in  ammonium  hydroxide.     Soluble  arsenates  precip.- 
tate    silver   (irtsntatc,    Ag3As04  ,    red-brown,    soluble    in    ammonium    hydroxide, 
nitric  acid,  arsenic  acid,  and  almost  insoluble  in  acetic  acid. 

A  solution  of  alkali  stannite — as  K2Sn02 — precipitates  metallic  silver 
from  solutions  of  silver  salts.  A  solution  of  silver  nitrate  in  a  great 
excess  of  ammonium  hydroxide  constitutes  a  very  delicate  reagent  to 
detect  the  presence  of  tin  in  the  stannous  condition  in  the  presence  of  fixed 
alkalis;  antimony  does  not  interfere  if  a  great  excess  of  ammonium  hy- 
droxide be  present. 

li. — Chromates  and  dichromates,  as  K2Cr04  and  K2Cr2O7  ,  precipitate  silver 
chromate,  Ag2Cr04  ,  dull-red,  sparingly  soluble  in  water  and  in  dilute  nitric 
acid,  soluble  in  ammonium  hydroxide. 

7.  Ignition. — Silver  nitrate  melts  undecomposed  at  218°,  at  a  red  heat  it  is 
decomposed  into  Ag°  ,  O  ,  N  ,  and  NO  (Fischer,  Pogg.,  1848,  74,  120).  Silver 
chloride  fuses  at  451°,  the  bromide  at  427°,  and  the  iodide  at  527°.  On  charcoal 
with  sodium  carbonate,  silver  is  reduced  from  all  its  compounds  by  the  blow- 
pipe, attested  by  a  bright  malleable  globule.  Lead  and  zinc,  and  elements  more 
volatile,  may  be  separated  from  silver  by  their  gradual  volatilization  under 
the  blow-pipe,  or  in  the  assay  furnace  (see  Cupellation  in  works  on  the  assay 
of  the  precious  metals). 


50  SILVER.  §59, 8. 

8.  Detection. — Silver  is  identified  by  its  precipitation  with  hydrochloric 
acid,  the  insolubility  of  the  precipitate  in  hot  water,  and  its  solubility  in 
ammonium  hydroxide,  with  reprecipitation  on  rendering  acid  with  nitric 
acid  (§61). 

9.  Estimation. — (a)  As  metallic  silver,  into  which  it  is  converted  by  direct 
ignition  if  it  is  the  oxide  or  carbonate,  or  by  ignition  in  hydrogen  if  the 
chloride,  bromide,  iodide  or  sulphide  (Vogel,  J.  C.,  1871,  24,  1009).  (b)  It  is 
precipitated  as  Ag<31  ,  and  after  igniting-  to  incipient  fusion,  weighed,  (c)  It  is 
converted  into  Ag,S  by  H2S  ,  and  weighed  after  drying-  at  100°;  inadmissible 
in  case  of  an  acid  that  might  liberate  free  sulphur,  (d)  Add  KCN  until  a 
solution  of  KAg(CN),  is  formed,  precipitate  with  HNO3  ,  and  after  drying  at 
100°,  weigh  as  AgCN  .  (e)  Volumetrically,  by  adding  a  graduated  solution  of 
NaCl  until  a  precipitate  is  no  longer  formed.  This  may  be  varied  by  adding 
the  measured  silver  solution  to  the  graduated  NaCl  solution,  containing  a  few 
drops  of  K2Cr04  ,  until  the  red  precipitate  begins  to  form,  (f)  Volumetrically, 
add  a  graduated  solution  of  ammonium  thiocyanate,  containing  ferric  sulphate, 
until  the  red  color  ceases  to  disappear,  (g)  Add  the  measured  silver  solution 
to  a  standard  solution  of  KCN  until  a  permanent  white  precipitate  is  formed. 

10.  Oxidation. — Metallic  silver  precipitates  gold  and  platinum  from 
their  solutions,  reduces  cupric  chloride  to  cuprous  chloride,1  mercuric 
chloride  to  mercurous  chloride,  and  permanganates  to  manganese  dioxide2. 
Silver  is  precipitated  from  its  solutions  by:  Pb  ,  PbS3.  Kg  ,  As4,  AsH3 , 
Sb ,  SbH3 ,  Sn ,  Sn",  Bi ,  Cu  ,  Cu'8,  Cd  ,  Te  ,  2'e  ,  FeS1,  Al ,  Mn  ,  Zn  ,  Mg  , 
P4,  PH3 ,  H,P02 ,  H2S03 ,  SiH4°,  H2026,  and  H  (very  slowly)7. 

In  alkaline  mixture  silver  is  also  reduced  by  Hg',  As"',  Sb'",  Bi'",  and 
Mn".  An  amalgam  of  mercury  and  tin  reduces  insoluble  compounds  of 
silver  in  the  wet  way,  the  silver  amalgamates  with  the  mercury  and  the 
tin  becomes  SnIV  (Laur,  C.  r.,  1882,  95,  38). 

Ferrous  sulphate  in  the  cold  incompletely  reduces  silver  salts;  on  boiling,  the 
ferric  salt  formed  is  reduced  and  the  silver  dissolved  (Lea,  7.  c.).  In  the  gradual 
reduction  of  silver  by  certain  organic  reagents,  the  metal  is  obtained  as  a  bright 
silver  coating  or  mirror  upon  the  inner  surface  of  the  test  tube  or  other  glass 
vessel.  Usually  a  slight^  ammoniacal  solution  of  silver  nitrate  is  used  and 
allowed  to  stand  some  time  with  the  reagent;  such  as  alcoholic  solution  of  oil 
of  cloves  or  cassia,  formic  acid,  aldehyde,  chloral,  tartaric  acid,  etc.  (ienlle 
warming  facilitates  the  result.  If  a  good  mirror  is  desired,  great  care  must  be 
taken  to  free  the  inner  surface  of  the  glass  from  all  organic  impurities  by 
careful  washing  with  ether,  chloroform,  etc.  In  these  deoxidations,  generally 
the  nitric  acid  radical  of  the  silver  nitrate  is  not  decomposed,  but  nitric  acid  is 
left:  4AgNO3  +  2HoO  =  4Ag-  -f  4HNO3  +  0,  . 

Light  acts  upon  nearly  all  salts  of  silver  when  mixed  with  gelatine  or  other 
organic  substances  used  in  preparing  photographic  plates,  etc.  These  plates 
contain  various  silver  salts,  frequently  the  bromide  or  iodide,  or  both  together. 
The  nature  of  the  chemical  change  is  not  fully  understood.  It  has  been  shown, 
however,  that  silver  chloride  on  exposure  to  the  light  loses  chlorine,  and  there 
is  considerable  evidence  to  prove  that  when  the  silver  halides  are  acted  upon 
by  light,  a  subhalide  such  as  Ag2Cl  ,  Ag2Br  or  Ag2I  is  formed.  When  the  pi  ate 
which  has  been  exposed  to  light  is  treated  with  a  reducing  agent,  the  reduc- 
tion of  the  silver  is  carried  to  the  metallic  state,  the  black  silver  producing  the 
image.  The  nitrate  in  crystal  or  pure  water  solution,  the  phosphate,  bromide, 

1  Lea,  Am.  S.,  1892,  144,  444.  "  /).,  2,  2,  759.  s  Skey,  C.  N.,  1871,  23,  232.  4  Senderens,  C.  r.,  1887, 
1O4.  175.  6  r>..  2. 1,  4515.  «  Rietfler,  J.  C.,  1896,  7O,  ii,  471.  7  Pellet,  B.,  1874,  7,  656  ;  Sehwarzenbach 

and  K  ritsvlK'u-sky,  ;..,  ISSU.  ar,,  1574  ;  <  'ooke,  C.  N..  1888.  58.  103.    «  Millon.  Am.  .S.,  18«i,  »«.  417. 


;60.  COMPARISON  OF  REACTIONS  OF  METALS  OF  THE  SILVER  GROUP.    51 


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§63,  60.  DIRECTIONS  FOE  ANALYSIS  WITH  NOTES.  53 

iodide  and  cyanide  are  not  decomposed  by  light  alone;  but  light  greatly  hastens 
their  decomposition  by  organic  substances,  or  other  reducing  agents,  as  of  solution 
of  silver  nitrate  in  rain  water,  or  \\ritten  as  an  ink  upon  fabrics.  Silver  is  the  base 
of  most  indelible  inks. 

DIRECTIONS  FOR  THE  ANALYSIS  OF  THE  METALS  OF  THE  FIRST  GROUP. 

§62.  Manipulation. — To  the  solution  acid  with  nitric  acid  add  hydro- 
chloric acid  (whenever  directions  call  for  the  addition  of  a  reagent  it 
is  to  be  used  reagent  strength  unless  otherwise  stated)  drop  by  drop 
(§32)  until  no  further  precipitate  is  formed  and  the  solution  is 
distinctly  acid  to  litmus  (§36).  The  precipitate  will  consist  of  the 
chlorides  of  Pb ,  Hgf,  and  Ag .  Shake  thoroughly  and  allow  to  stand 
a  few  moments  before  filtering;  if  the  solution  is  warm  it  should 
be  cooled  to  the  temperature  of  the  room.  Decant  the  solution  and 
precipitate  upon  a  filter  paper  previously  wetted  (§35)  with  water  and 
wash  two  or  three  times  with  cold  water  or  until  the  filtrate  is  not  strongly 
acid  to  litmus.  The  washings  with  cold  water  should  be  added  to  the 
first  filtrate  and  the  whole  marked  and  set  aside  to  be  tested  for  the 
metals  of  the  remaining  groups  (§16). 

§63.  Notes. — 1.  Failure  to  obtain  a  precipitate  upon  the  addition  of  HC1  to 
an  acid  reaction  is  proof  of  the  absence  of  Hg'  and  Ag ,  but  a  solution  of  a 
lead  salt  may  be  present,  of  such  a  degree  of  dilution  that  the  lead  chloride 
formed  will  be  soluble  in  the  dilute  acid  (§57,  5c). 

2.  The   solution   should   not  be   strongly   acid  with   nitric   acid,   as  it   forms 
nitrohydrochloric   acid  with   the   hydrochloric   acid,   causing  oxidation   of  the 
Hg'  (§58,  5c).     Lead  chloride  is  also  more  soluble  in  nitric  acid  than  in  dilute 
hydrochloric  acid  (§57,  5c).     By  a  study  of  the  solubilities  of  the  silver  group 
metals  it  will  be  seen  that  H2S04  ,  HC1  ,  HBr  or  HI  cannot  be  used  in  prepar- 
ing- a  solution  for  analysis  when  these  metals  are  present. 

3.  A  great  excess  of  acid  is  to  be  avoided,  as  it  may  interfere  with  the  reac- 
tion  in  Group   II.    (§57,    6e).      Complete    precipitation    should   be    assured    by 
testing  the  filtrate  with  a  drop  of  HC1  ,  when  no  further  precipitation  should 
occur   (§32).     If  a  white  precipitate  is  formed  by   adding  a  drop   of  HC1  to 
the  filtrate  it  is   evident  that  the   precipitation   was  not   complete   and   more 
HC1  should  be  added  and  the  group  separation  repeated  (§11). 

4.  The  presence  of  a  slight  excess  of  •dilute  acid  does  not  aid  or  hinder  the 
precipitation   of  the  Hg'  or  Ag ,  but   as  PbCl2  is   less   soluble   in  dilute   HC1 
than  in  water,  a  moderate  excess  of  the  acid  causes  a  more  complete  precipita- 
tion of  that  metal  in  the  first  group. 

5.  Concentrated  HC1  dissolves  the   chlorides   of  the  first   group   quite  appre- 
ciably (§59,  5c). 

6.  Hydrochloric    acid    added   to    certain    solutions    may    cause    a    precipitate 
when  none  of  the  first  group  metals  are  present.     Some  of  the  more  important 
conditions  are  mentioned.     It  will  be  noted  that  a  number  of  these  are  alkaline 
solutions  and  will  give  precipitates  with  other  acids  than  HC1.     It  is  advisable 
in  these  cases  to  acidify  with  HNO3  before  adding  HC1  in  order  to  avoid   error 
from  this  source. 

a.  A  concentrated  solution  of  BaCL  is  precipitated  without  change  by  the 
addition  of  HC1 ,  readily  soluble  in  water  (§186,  5c). 

b.  An   acid   solution  of   Sb  ,   Bi ,   or    Sn  ,   with   some   other   acid   than   HC1  , 
and   saturated   with   water  as   far   as   possible   without   precipitation,    on    the 
addition    of    HC1  ,    precipitates    the    oxychloride    of    the    corresponding    metal 
(§76,  G/1).     These  precipitates  are  readily  soluble  in  an  excess  of  the  HC1  .     It 
must,  however,  be  remembered  that  a  trace  of  AgCl  will  also  be  dissolved  by 
an  excess  of  HC1  (§59,  5c). 


54  DIRECTIONS  FOR  ANALYSIS   WITH  NOTES.  §63, 6/. 

c.  Solutions  of  metallic  oxides  in  the  alkali  hydroxides  are  precipitated  when 
neutralized  with  acids,  e.  g.,  K2ZnO2  +  2HC1  =  Zn(OH)2  +  2KC1 . 

d.  r\  he  sulphides  of  As ,  Sb ,  Sn ,  Au ,  Pt ,  Mo  (Ir ,  W ,  Ge  ,  V ,  Se  and  Te) 
in    solution    with    the    alkali    polysulphides    are     reprecipitated     together    with 
sulphur  on  the  addition  of  HC1  (§69,  6e). 

e.  Soluble    polysulphides    and    thiosulphates    give    a    precipitate    of    sulphur, 
white,  with  HC1  (§256,  3o). 

f.  Certain    soluble    double    cyanides,    as    Ni(CN)3.2KCN'  ,    are    precipitated 
as  insoluble  cyanides,  Ni(CN)2  ,  on  the  addition  of  HC1  (§133,  (>7>). 

g-.  Solutions  of  silicates  (§249,  4),  borates,  tuiigstates,  molybdates;  also 
benzoates,  salicylates,  urates,  and  certain  other  organic  salts,  are  precipitated 
by  acidulation  with  HC1 ,  many  of  the  precipitates  being-  soluble  011  further 
addition  of  the  acid. 

h.  Acidulation  with  HC1  may  induce  changes  of  oxidation  or  reduction, 
which  in  certain  mixtures  may  result  in  precipitation:  for  example,  Cu"  salts 
with  KCNS  in  ammoniacal  solution  (§77,  M) ;  mixture  of  solutions  of  KI  and 
KI03  (§280,  6,5,7),  etc. 

7.  If   the   precipitate,   obtained   by   the  addition   of   HC1  to   the   solution,  is 
colored  or  does  not  give  further  reactions  which  are  conclusive  and  perfectly 
satisfactory  in  every  respect,  it  should  be  separated'  by  filtration,  and  treated 
as   a  solid    substance   taken    for    examination    (see    conversion    of    solids    into 
liquids,  §301). 

8.  Compounds  of  the  first  group  metals  insoluble  in  water  or  acids  are  trans- 
posed to  sulphides  by  digestion  with  an  alkali  sulphide.     The  lead  and  silver 
sulphides  thus  formed  are  readily  soluble  in  hot  dilute  nitric  acid.     The  mer- 
curous  compounds  are  changed  to  mercuric  sulphide  (§58,  5tt  and  (te),  a  second 
group  mercury  compound  insoluble  in  HNO3  . 

9.  If  but  one  metal   of  the  first  group  be   present,   the   action   of   NH4OH 
determines  which  it  is;  PbCl2  does  not  change  color  or  dissolve;  HgCl  blackens; 
and  AgCl  dissolves  (§60). 

§64.  Manipulation. — The  precipitate  (white)  on  the  filter  should  now 
be  washed  once  or  twice  with  hot  water.  The  first  hot  water  should  be 
poured  upon  the  precipitate  a  second  time.  This  hot  filtrate  is  divided 
into  four  portions  and  each  portion  tested  separately  for  lead  with  the 
following  reagents,  H2S04 ,  H2S ,  K2Cr207 ,  and  KI  (§57,  6  e,  h,  and  /) 
giving  white,  black,  and  yellow  precipitates : 

The  yellow  precipitate  with  potassium  iodide  (the  KI  must  not  be  used 
in  great  excess  (§57,  5c))  should  he  allowed  to  settle,  the  liquid  decanted, 
and  the  precipitate  redissolved  in  hot  water,  to  a  colorless  solution  which 
upon  cooling  deposits  beautiful  yellow  crystalline  scales  of  PbI2  (charac- 
teristic of  lead). 

§65.  Notes. — 1.  Lead  is  never  completely  precipitated  in  the  first  group 
(§57,  Of).  The  presence  of  a  moderate  excess  of  dilute  HC1  and  the  cooling  of 
the  solution  both  favor  the  precipitation. 

2.  Lead  can  be  completely  separated  from  the  second  group  metals  by  sul- 
phuric acid  applied  to  the  original  solution   (§57,  6e,  §95  and  §98),  but  that 
would  necessitate  a  regrouping  of  the  metals;  as,  Ba  ,  Sr  ,  and  Ca  would  also 
be  precipitated  (Zettnow,  Z.,  1867,  6,  438). 

3.  In  order  to  precipitate  the  lead  chloride,  not  removed  in  the  first  group,  in 
the  second  group  with  H2S  ,  the  solutions  must  not  be  strongly  acid,  either 
the  excess  of  HC1  should  be  removed  by  evaporation  or  the  solution  should  be 
diluted  (§57,  6e,  and  §81,  3,  5  and  9). 

4.  If  the  lead  chloride  is  not  all  washed  out  with  hot  water  it  is  changed  to 
an  insoluble  basic  salt  (white)  by  the  NH4OH  ,  part  remaining  on  the  filter 
and   part   carried   through   mechanically  which   causes   turbidity   to    the    am- 


§68,  3.  DIRECTIONS  FOR  ANALYSIS   WITH  NOTES.  55 

monium  hydroxide  solution  of  the  AgCl  and  makes  necessary  the  filtration  of  that 
solution  before  the  addition  of  HNO3 ,  otherwise  it  does  not  interfere. 

5.  The  precipitation  of  lead  as  the  sulphide  while  not  characteristic  of  lead, 
is  exceedingly  delicate,  much  more  so  than  the  formation  of  the  white  PbSO4 
(§57,  5c).     In  extremely  dilute  solutions  no  precipitate  occurs,  merely  a  brown 
coloration  to  the  solution.     The  presence  of  free  acid  lessens  the  delicacy  of 
the  test. 

6.  PbCrO4  is  blackened  by  alkali  sulphides  and  dissolved  by  the  fixed  alkalis 
(important  dislinclioii  from  BaCr04);  the  solubility  in  the  fixed  alkalis  is  also 
an  important  distinction  from  bismuth  chromate  (§76,  Gh). 

7.  Other  tests  for  lead  by  reduction  on  charcoal  before  the  blow-pipe,  or  in 
the  wet  way  by  Zn,  should  not  be  omitted  (§57,  7  and  10).     If  to  a  solution  of 
lead  salt  nearly  neutral  a  strip  of  zinc  be  added,  the  lead  will  soon  be  deposited 
on  the  zinc  as  a  spongy  mass. 

§66.  Manipulation. — The  white  precipitate  remaining  on  the  filter  after 
washing  with  hot  water  consists  of  HgCl  and  AgCl ,  and  some  PbCl2  if  it 
has  not  been  well  washed.  To  this  precipitate  NH^OH ,  one  or  two  cc. 
is  added  and  allowed  to  pass  through  the  filter  into  a  clean  test-tube.  An 
instantaneous  blackening  of  the  precipitate  is  evidence  of  the  presence  of 
mercurous  mercury. 

The  AgCl  is  dissolved  by  the  NH4OH ,  and  is  found  in  the  filtrate ;  its 
presence  being  confirmed  by  its  reprecipitation  on  rendering  the  solution 
acid  with  HN03 . 

§67.  Notes. — Mercury. — 1.  The  black  precipitate  on  the  filter,  caused  by  the 
addition  of  NH4OH  to  the  HgCl  may  be  examined  under  the  microscope  for 
the  detection  of  globules  of  Hg°,  or  the  precipitate  may  be  digested  with  con- 
centrated solution  of  (NH4)2  SO4,  which  dissolves  the  NH2HgCl ,  leaving  the 
Hg°  unattacked  (§58,  60). 

2.  If    the    original    solution    contains    no    interfering    metals,    the    distinctive 
reactions   of   mercurous   salts   with   iodides,    chromates   and   phosphates   should 
be  obtained  (§58,  6e,  h  and  d). 

3.  Mercury    has  but  few  soluble   mercurous    compounds,    and    in    preparing 
solutions  of  the  insoluble  compounds  for  analysis,  oxidizing  agents  are  usually 
employed  and  the  mercury  is  then  found  entirely  in  the  second  group  as   a 
sulphide  (§96  and  §97). 

4.  Additional  proof  may  be  obtained  by  mixing  a  portion  of  the  black  residue 
with  sodium  carbonate,  drying  and  heating  in  a  glass  tube  (read  §58,  7,  also 
§97,  7). 

§68.  Silver. — 1.  The  presence  of  a  large  excess  of  Hg(N03)2  prevents  the 
precipitation  of  AgCl  from  solutions  of  silver  salts  by  HC1  (§59,  5<0-  In  this 
case  the  metals  should  be  precipitated  by  H,S  and  the  well-washed  precipitate 
digested  with  hot  dilute  HNO3  .  The  silver  "is  dissolved  as  AgNO,,  .  while  the 
mercury  is  unattacked:  6Ag,S  -f  16HNO8  =  12AgNO3  +  3S2  +  4NO  +  8H2O  . 
After  evaporation  of  the  excess  of  HNO3  the  solution  may  be  treated  with 
HC1  as  an  original  solution. 

.  A  small  amount  of  AgCl  with  a  large  amount  of  HgCl  is  not  dissolved  by 
NH4OH  ,  but  is  reduced  to  Ag°  by  the  Hg°  formed  bv  the  addition  of  the 
1TH4OH  to  the  HgCl  (§58,  0</,  §59.  6/,  10  and  §60). 

3.  If  Hg'  be  present  and  Ag-  is  not  detected  the  black  precipitate  on  the 
filter  should  be  digested  for  some  time  with  (NH4)2S  ,  washed,  and  boiled  with 
hot  dilute  nitric  acid.  The  Ag  ,  if  any  be  present,  is  dissolved  and  separated 
from  the  HgS : 

NH.HgCl  +  (NH4),S  +  2H20  =  HgS  -f  NH4C1  -f  2NH.OH 

Hg  +  (NH4)aSx  =  HgS  +  (ira4),S..l 


56  ARSENIC.  §68,  4. 

The  silver  may  also  be  detected  by  treating  the  precipitate  of  HgCl  and 
AgCl  with  an  acid  oxidizing  agent  such  as  HC1  and  KC1O3.  The  HgCl  will  be 
oxidized  to  HgCl2  and  dissolved  while  the  AgCl  remains  undissolved. 

The  silver  may  also  be  detected  as  follows:  Boil  the  original  solution  with 
K2S2O4  until  the  Hg'  is  oxidized  to  Hg"  then  add  one  drop  of  MnSO,  and  more 
K2S2O4  and  boil,  if  the  manganese  is  oxidized  to  HMnO4  it  proves  that  silver 
is  present. 

4.  If  only  a  trace  of  silver  be  present,  its  detection  by  adding-  HN03  to  the 
NH4OH  solution  of  the  chloride  may  fail,  unless  the  excess  of  the  NH4OH  be 
first  removed   by  evaporation    (because   of  the   solubility   of  the  AgCl  in   the 
ammonium  salt,  §59,  5e). 

5.  As  a  further  test  for  silver,  the  chloride,  precipitated  by  the  nitric  acid, 
may  be  reduced  to  the  metal  by  zinc;  by  adding-  to  the  ammoniacal  solution 
a  few  drops  of  potassium  stannite   (§71,  6a  and  8);  by  warming-  with  grape 
sugar  in  alkaline  mixture.     In  all  cases  the  well-washed  grayish  black  metal 
may  be  dissolved  in  nitric  acid  as  Ag-N03  . 

6.  To  identify  the  acid  of  silver  salts  which  are  insoluble  in  HN03(AgCl, 
AgBr ,  Agl),  (1)  Add  metallic  zinc  and  a  drop  of  H2SO4 ;  when  the  silver  is  all 
reduced  test  for  the  acid  in  the  filtrate.     (2)  Fuse  with  Na2C03  ,  add  water, 
and  test  the  filtrate   for  acids.     (3)    Add  HL,S  ,   or   an    alkali   sulphide,   digest 
warm  for  a  few  minutes,  filter  and  test  filtrate  for  acids.     (.'/)  Boil  with  KOH 
or  NaOH  (free  from  HC1),  and  test  the  filtrate  in  the  same  manner.     It  must 
not   be   overlooked   that   by    the    first   three    methods,    and    not    by    the    last, 
bromates  and  iodates  are  reduced  to  bromides  and  iodides  (§257,  62?). 


THE  TIN  AND  COPPER  GROUP  (SECOND  GROUP). 

Arsenic,  Antimony,  Tin,  Gold,  Platinum,  Molybdenum,  Mercury,  Lead, 
Bismuth,  Copper,  Cadmium  (Buthenium,  Khodium,  Palladium,  Iridium, 
Osmium,  Tungsten,  Vanadium,  Germanium,  Tellurium,  Selenium). 

THE  TIN  GROUP  (SECOND  GROUP,  DIVISION  A). 

Arsenic,  Antimony,  Tin,  Gold,  Platinum,  Molybdenum  (Iridium,  Tungs- 
ten, Vanadium,  Germanium,  Selenium,  Tellurium). 

§69.  Arsenic.    As  =  74.96.     Valence  three  and  five. 

1.  Properties. — Specific  gravity,  pure  crystalline  5.727  at  14°;  amorphous 
4.716  (Bettendorff,  A.,  1867,  144,  110).  Melting  point,  500°  (Burgess,  Wash. 
Acad.  of  Sc.,  1-18);  between  the  melting  point  of  antimony  and  silver  (Mallet, 
C.  N.,  1872,  26,  97).  Volatilizes  in  an  atmosphere  of  coal  gas  without  melting 
at  450°  (Conechy,  C.  N.  1880  41,  189).  Vapor  aensity  (H  =  1),  147.2  (Deville 
and  Froost  C.  r.,  1863,  66,  891);  therefore  the  molecule  is  assumed  to  con- 
tain four  atoms  (As4).  At  a  white  heat  the  vapor  densty  is  less,  but  the  dis- 
sociation is  not  low  enough  to  indicate  As2  (Mensching  and  V.  Meyer,  B.,  1887, 
20,  1833).  Arsenic  exists  in  two  forms,  crystalline  and  amorphous.  The  crys- 
talline arsenic  is  steel-gray  with  a  metallic  luster,  brittle  and  easily  pulyerizable; 
forms  beautiful  rhombic  crystals  on  sublimation  with  slow  condensation.  For 
ductility,  malleability,  etc.,  see  D.,  2,  1,  161.  Amorphous  arsenic  is  grayish 
black,  of  less  specific  gravity  than  the  crystalline;  long  heating  changes  it  to 
the  crystalline  form  (Engel,  C,  r.,  1883,  96,  1314).  The  vapor  of  arsenic  is  citron- 
yellow  (Le  Roux,  C.  r.,  1860,  61,  171),  with  an  oppressive  and  poisonous  alli- 
aceous odor.  It  is  slowly  oxidized  in  moist  (not  in  dry)  air  at  ordinary  temper- 
ature; when  heated  in  the  air,  it  burns  with  a  bluish  flame  and  becomes  the 
white  arsenous  anhydride,  As2O3  .  The  burning  metal  evolves  a  strong  garlic 
odor,  not  noticed  when  the  pure  arsenous  anhydride  is  sublimed.  In  its  phys- 
ical properties  arsenic  is  a  metal,  but  its  failure  to  act  as  a  base  with  oxyacids 
classes  it  chemically  with  the  non-metallic  elements  (Adie,  J.  C.,  1889,  55, 
157;  Stavenhagen,  Z.  angew.,  1893,  283).  Its  chief  use  as  a  metal  is  in  mixing 
with  lead  for  making  shot. 


§69,  5b.  ARSENIC.  57 

2.  Occurrence. — Arstenic    is    very    widely    distributed    geographically.     It  is 
occasionally  found  native;    the  chief  arsenic  minerals  are  realgar,  AsS,  and  orpi- 
ment,    As2S3  ;     it   is    an    essential    constituent    of    many    other    minerals,     e.    </., 
niccolite,  NiAs  ;    cobaltite,   CoAsS  ;    smaltite,    (Cp,Ni)As2  ;  arsenopyrite,  FeAsS  ; 
lollingite,     FeAs2  ;      proustite,     Ag3AsS3  ;      enargite,     Cu3AsS4  ;      erythrite,     Co3 
(AsO4)28H2O  ,     and     many     others.       Many    sulphide    ores    of    zinc    and    iron 
contain   arsenic,   hence  arsenic  is  frequently  found  in  these   metals  and  in  sul- 
phuric acid  made  from  the  sulphur,'  and  also  in  the  products  made  therefrom. 

3.  Preparation. — (1)  Reduced  from  its  oxide  by  ignition  with  carbon;  JAs  O 
+  3C  =  As4  +  3CO2  .     (2)   From  arsenopyrite,   FeAsS  ,   by  simple  ignition,   air 
being  excluded;    4FeAsS  =  4FeS  +  As4  .     (3)     From  orpiment,  As2S3  ,  by  fusion 
with    sodium    carbonate    and    potassium    cyanide;      2As2Ss  +  6Na2CO3  +  6KCN 
=  Aa<  +  6Na2S  +  6KCNO  +  6CO2  . 

4.  Oxides. — Arsenic  forms  two  oxides:  arsenous  oxide  or  anhydride,  As,,03 
(Biltz,  Z.  phys.  Ch.,  1896,  19,  385;  C.  C.,  1896,  793),  and  arsenic  oxide  or  anhydride, 
As2O5  .    .Arsenous  oxide,  As,O3  (white  arsenic,  arsenous  anhydride,  arsenous  aoid, 
arsenic  trioxide),  is  usually  prepared  by  burning  arsenic;  it  may  also  be  prepared 
by  heating-  arsenic  in  sulphuric  acid   till   S02   is   evolved,   or  by   decomposing 
AsCl3  with  H2O  .     It  sublimes  easily  on  gradually  heating,  forming  beautiful 
octahedral  and  tetrahedral  crystals.     On  suddenly  heating  under  pressure  it 
melts,  and  on  cooling  forms  the  opaque  arsenic  glass.     It  is  very  poisonous, 
usually   producing  violent   vomiting.     One   hundred  fifty   milligrams   are   con- 
sidered a  fatal  dose  for  an  adult.     No  acids  (hydroxides)  of  arsenous  anhydride 
(oxide)  have  been  isolated;  but    it  reacts  with  bases,   forming  salts,   arsenites, 
as   if    derived    from   the   meta,   ortho,    and    pyro   arsenous    acids.     The    alkali 
arsenites  are  usually  meta  compounds;  the  arsenites  of  the  alkaline  earths  and 
heavy  metals  are  usually  ortho  compounds  (D.,  2,  1,  170). 

Arsenic  pentoxide,  As2O5  (arsenic  anhydride,  arsenic  oxide),  is  formed  by  heat- 
ing arsenic  acid,  H3As04  (Berzelius,  A.  Ch.,  1819,  11,  225).  It  is  a  white 
amorphous  mass,  melts  at  a  dull  red  heat,  is  slowly  deliquescent,  combining 
with  water  to  form  H8As04  .  The  pentoxide,  As2O5  ,  forms  three  acids  or 
hydroxides:  meta-arsenic  acid,  HAsO8  =  AsO2(OH);  ortho-arsenic  acid, 
H3As04  =  AsO(OH)3;  and  pyro-arsenic  acid,  H4As2O7  —  As203(OH)4;  each 
of  these  forming  a  distinct  class  of  arsenates  with  bases.  Ortho-arsenic  acid  is 
formed  by  adding  water  to  arsenic  anhydride,  As,O ,  -f  3H2O  =  2H3AsO4  , 
or  by  oxidizing  arsenic  or  arsenous  anhydride  with  nitric  acid.  Pyro-arsenic 
acid  is  formed  by  heating  the  ortho  acid  to  between  140°  and  180°:  2H3As04  = 
H4As2O7  -f-  H2O  .  The  meta  acid  is  formed  by  heating  the  ortho  or  pyro  acid 
to  206°:  H3AsO4  =  HAs03  +  H2O  (D.,  I.  c.). 

5.  Solubilities. — a. — Metal. — Arsenic  is  insoluble  in  pure  water.     It  is  readily 
attacked  by  dry  chlorine  and  bromine  upon  contact  and  by  iodine  with  the  aid 
of   heat.     Arsenous    chloride,    bromide   and   iodide    are    formed.     It    combines 
with  sulphur,  forming  from  As,S2  to  As2S5  ,  depending  upon  the  proportion  of 
sulphur  present   (Gelis,  A.  Ch.,   1873,   (4),  30,  114).     Chlorine  and  bromine  in 
presence  of  water  oxidize  it,  first  to  arsenous  then  to  arsenic  acid    (Millon, 
A.  Ch.,  1842,  (3),  6,  101):  As4  -f  10C12  +  16H2O  =  4H3AsO4  +  20HC1  .     It  is  not 
attacked  by  concentrated  hydrochloric  acid  at  ordinary  temperature  and  but 
slowly  by  the  hot  acid  in  presence  of  air  forming  As203  ,  then  AsCl3  ;  nitric 
acid  readily  oxidizes  it  first  to  As203  then  to  H3AsO4  ;  upon  fusion  with  KNO, 
it    becomes    K3AsO4;    readily    soluble    as   H3As04    by    nitrohydrochloric    acid; 
sulphuric  acid,  dilute  and  cold,  is  without  action;  with  heat  and  the  more  con- 
centrated  acid   As203    is   formed   and   the   sulphuric   acid   is   reduced   to   SO3  . 
Ammonium  hydroxide  is  without  action   (Guenez,  C.  r.,  1892,   114,  1186).     Hot 
solution   of   potassium   or  sodium   hydroxide   dissolves   it   as   arsenite:   As4    + 
4KOH  -f  4H2O  =  4KAsO2  -f-  6H2  . 

1>- — Oxides. — Arscnous  oxide  exists  in  two  forme,  crystalline  and  amorphous,  the 
solubilities  of  which  differ  considerably  (§27).  At  ordinary  temperature  100 
parts  of  water  dissolve  3.7  parts  of  the  amorphous  and  1.7  parts  of  the  crystal- 
line, several  hours  being  necessary  to  effect  the  solution.  100  parts  of  boiling 
water  dissolve  11.46  parts  of  the  amorphous  and  10.14  parts  of  the  crystalline 
oxide  in  three  hours  (Winkler,  J.  pr.,  1885,  (2),  31,  247).  The  presence  of  acids 
greatly  increases  the  solubility  in  water  (Schultz-Sellac,  B.,  1871,  4,  109). 
Arsenous  oxide  is  readily  soluble  in  alkali  hydroxides  or  carbonates  to  arsenites 


58  ARSENIC.  §69, 5c. 

(Clayton,  C.  N.,  1891,  64,  27),  Arsenic  'pcnto,ri(lc,  As2On  ,  is  deliquescent,  soluble 
in  water  forming  H3As04  .  The  meta  and  pyro  acids  are  easily  soluble  in 
water  forming  the  ortho  acid  (Kopp,  A.  Ch.,  1856,  (3),  48,  106). 

c. — Salts. — Arsenic  does  not  act  as  a  base  with  oxyacids,  but  its  oxides  combine 
with  the  metallic  oxides  to  form  twro  classes  of  salts,  arsenites  and  arsenates. 
Arsenites  of  the  alkalis  are  soluble  in  water,  all  others  are  insoluble  or  only 
partially  so;  all  are  easily  soluble  in  acids.  Alkali  arsenates,  and  acid  arsenates 
of  the  alkaline  earths,  are  soluble  in  water;  all  are  soluble  in  mineral  acids, 
including  H3AsO4  (LeFevre,  C.  r.,  1889,  108,  1058).  See  also  under  the  respec- 
tive metals. 

Arsenous  sulphide,  As2S3 ,  is  insoluble  in  water  when  prepared  in  the  dry 
way ;  when  prepared  in  the  moist  way  it  may  be  transformed  into  the  soluble 
colloidal  *  form  by  treatment  with  pure  water,  from  which  solutions  it  is 
precipitated  by  solutions  of  most  inorganic  salts  or  acids  (Schulze,  J.  pr., 
1882  (2),  25,  431).  The  presence  of  acids  or  solutions  of  salts  prevents 
the  solubility  of  As2S3  in  water.  Boiling  water  slowly  decomposes  the 
sulphide  forming  As203  and  H2S  (Field,  C.  N.,  1861,  3,  115.;  Wand,  Arch. 
Phar.,  1873,  203,  296).  It  is  completely  decomposed  by  gaseous  HC1  form- 
ing AsCl3  (Piloty  and  Stock,  R.,  1897,  30,  1649),  very  slightly  decomposed 
and  the  arsenic  dissolved  by  hot  concentrated  acid  (Field,  I.  c.).  Chlorine 
water  and  nitric  acid  decompose  it  readily  with  formation  of  H3As04; 
with  sulphuric  acid  As203  and  S02  are  formed  (Rose,  Pogg.,  1837,  42,  536). 
The  alkali  hydroxides  or  carbonates  dissolve  it  readily  with  formation  of 
RAs02  and  EAsS2  (E  =  K ,  Na  and  NHJ  V.,  2,  1,  183) ;  soluble  in  alkali 
sulphides  and  poly-sulphides  forming  R4As2S5 ,  and  RAsS2  (Berzelius, 
Pogff.-,  1826,  7, 137 ;  Nilsson,  J.  0. ,  1872,  25,  599).  Whether  the  ortho,  meta 
or  pyro  salt  is  formed,  depends  upon  the  amount  of  alkali  sulphide  present 

Arsenic  sulphide,  As2S5 ,  is  insoluble  in  water ;  soluble  in  HC1  gas,  as 
AsCl3  ;  insoluble  in  dilute  HC1 ,  soluble  in  HN03  or  chlorine  water,  as 
H3As04  ;  soluble  in  alkali  hydroxides  and  carbonates,  as  R3AsS4  and 
R3As03S  :  As2S5  -f  6NH4OH  =  (NH4)3AsS4  +  (NH4)3As03S  +  3H20  (Mc- 
Cay,  Ch.  Z.,  1891,  15,  476);  soluble  in  alkali  sulphides,  as  K3AsS4  (Nilsson, 
J.  pr.,  1876  (2),  14,  171). 

Arsenous  chloride,  bromide  and  iodide  (AsCl3  ,  AsBr,  ,  Asl,)  are  decomposed 
by  small  amounts  of  water  into  the  corresponding  oxyhalogen  compounds, 
AsOCl  ,  etc.  A  further  addition  of  water  decomposes  these  compounds  into 
arsenous  oxide  and  the  halogen  acids. 

6.  Reactions. — a. — The  alkali  hydroxides  and  carbonates  unite  with  arsenous 
and  arsenic  oxides  (acids),  the  latter  with  evolution  of  carbon  dioxide,  forming- 
soluble  alkali  arsenites  and  arsenates.  These  alkali  salts  are  chiefly  meta  arse- 
nites and  ortho  arsenates  (Bloxam,  J.  (7.,  1862,  15,  281;  Graham,  Pogff.,  1834,  32, 
47). 

*  Colloids  is  a  name  given  by  Graham  to  a  class  of  glue-like  bodies  in  distinction  to  the  crystal- 
loids,  which  have  a  weU-defined  solid  form.  The  colloids  are  indefinitely  soluble  in  water, 
giving  the  little-understood  "  pseudo-solutions,"  which  stand  midway  between  the  mechanical 
suspension  or  emulsion  and  the  true  solution.  Gelatine,  starch,  the  metallic  sulphides,  silicic 
acid,  and  the  hydroxides  of  iron  and  aluminum  are  some  of  the  substances  that  may  take  on  the 
colloid  torn,  TBe  colloid  lolutioni  are  ai  a  rule  broken  up  by  addition  of  an  »ei4  or  a  neutral 
H4t, 


§69,  6e.  ARSENIC.  59 

ft.  —  Oxalic  acid  does  not  reduce  arsenic  acid*  (Nay  lor  and  Braithwaite,  Pharm. 
./.  Trans.,  1883,  (3),  13,  464).  Potassium  ferricyaiiide  in  alkaline  solution  oxi- 
dizes arsenous  compounds  to  arsenic  compounds,  very  rapidly  when  gently 
warmed,  c.  Nitric  acid  readily  oxidizes  all  other  compounds  of  arsenic  to 
arsenic  acid.  d.  Hypophosphites  in  presence  of  concentrated  hydrochloric  acid 
reduce  all  oxycompounds  of  arsenic  to  the  metallic  state.  0.00001  gram  oi 
arsenic  may  be  detected  by  boiling  with  10  cc.  strong  hydrochloric  acid  and  0.2 
gram  calcium  hypophosphite  (Engel  and  Bernard,  C.  r.,  1896,  122,  390;  Thiele 
and  Loof,  C.  C.,  1890,  1,  877  and  1078;  and  Hager,  J.  C.,  1874,  27,  868). 

e.  —  Hydrosulphuric  acid  precipitates  the  lemon-yellow  arsenous  sulphide, 
As2S3  ,  from  acidulated  solutions  of  arsenous  acid.  The  precipitate  forms 
in  presence  of  concentrated  hydrochloric  acid.  Citric  acid  and  other 
organic  compounds  hinder  the  formation  of  the  precipitate,  but  do  not 
wholly  prevent  it  if  strong  hydrochloric  acid  be  present.  Nitric  acid 
should  not  be  present  in  strong  excess  as  it  decomposes  hydrosulphuric 
acid,  with  precipitation  of  sulphur. 

In  aqueous  solutions  of  arsenous  acid  the  sulphide  forms  more  as  a 
yellow  color  than  as  a  precipitate,  being  soluble  to  quite  an  extent  in  pure 
water,  especially  when  boiled  (5c)  :  As,S3  +  3H,0  =  As203  +  3H2S  .  This 
has  been  given  as  a  method  of  separating  arsenous  sulphide  from  all  other 
heavy  metal  sulphides  (Clermont  and  Frommel,  /.  C.,  1879,  36,  13).  The 
precipitate  is  not  formed  in  solutions  of  the  arsenites  except  upon  acidu- 
lation.  The  hydrogen  sulphide  converts  the  oxy  salts  of  arsenic  into  the 
thio  salt,  which  is  decomposed  by  acid  with  precipitation  of  the  sulphide 
of  arsenic  : 

Na.iAsOa  +  3H2S  =  Na3AsS3  +  3H2O 
2Na3AsS3  +  6HC1  =  As2S3  +  3H2S 

Alkali  sulphides  produce  and,  by  further  addition,  dissolve  the  precipi- 
tate (5c)  : 

2AsCl3  +  3(NH4)2S  =  AS2S3  +  NH4C1 

As2S3  +  2(NH4)S  =  (NH4)4AS2S5  or  A2S3  +  (NH4)2S  =  2NH4AsS2 

Arsenous  sulphide  is  also  soluble  in  alkali  hydroxides  and  carbonates,  with 
evolution  of  C02,  forming  arsenites  and  thioarsenites  (5c).     The  thioarse- 
nites  are  precipitated  by  acids  forming  As2S3   :  (NH4)  4^8085  +  4HC1  = 
As2S3  +  2H2S  +  4NH4C1  or  2NH4AsSt  +  2HC1  =  As2S3  +  H2S    + 


The  solubility  of  the  sulphides  of  arsenic  in  yellow  ammonium  sulphide 
separates  arsenic  with  antimony  and  tin  from  the  other  more  common 
metals  of  the  second  group;  and  the  solubility  in  ammonium  carbonate 
effects  an  approximate  separation  from  antimony  and  tin  (Hager,  J.  C., 
1885,  48,  838).  Arsenous  sulphide  is  soluble  in  solutions  of  alkali  sul- 
phites containing  free  sulphurous  acid  (separation  from  antimony  and 
tin):  4As2S3  +  32KHSO,,  =  8KAs02  +  12K,S,0,  +  3S2  +  US02  +  16H20. 
It  may  also  be  separated  from  antimony  and  tin  by  boiling  with  strong 
hydrochloric  acid,  the  As2S.<  remaining  practically  insoluble;  the  sulphides 
of  antimony  and  tin  being  dissolved.  It  is  easily  dissolved  by  strong 


60  AltsmiC.  §69,  6/1 


hydrochloric  acid,  the  As  S;  remaining  practically  insoluble;  the  sulphides 
of  antimony  and  tin  being  dissolved.  It  is  easily  dissolved  by  strong 
nitric  acid,  and  by  free  chlorine  or  nitrohydrochloric  acid,  as  arsenic  acid: 
6As2S3  +  20HN03  +  SH20  =  12H3As04  +  9S2  +  20NO  ;  2As2S3  +  10CL, 
+  16H20  =3  4H3As04  +  3S2  +  20HC1  .  Usually  a  portion  of  the  sulphur 
is  oxidized  to  sulphuric  acid,  completely  if  the  nitric  acid  or  chlorine  be  in 
great  excess  and  heat  be  applied:  As2S3  +  14C12  +  20H20  =  2H3As04  + 
3H2S04  +  28HC1  . 

Arsenic  pentasulpliide,  As2S5  ,  is  formed  by  passing  H2S  for  a  long  time 
into  a  solution  of  alkali  arsenate  and  then  adding  acid  (McCay,  Am.,  1891, 
12,  547);  by  saturating  a  solution  of  arsenic  acid  with  H2S  and  placing,  in 
stoppered  bottle,  in  boiling  water  for  one  hour;  or  by  passing  a  rapid 
stream  of  H2S  into  an  HC1  solution  of  H3As04  (Bunsen,  A.,  1878,  192,  305  ; 
Brauner  and  Tomicek,  J.  C.,  1888,  53,  146);  2H3As04  +  5H2S  +  xHCl  = 
As,S.  -f-  8H20  -f-  xHCl  .  Carbon  disulphide  extracts  no  sulphur  from  the 
precipitate,  indicating  the  absence  of  free  sulphur.  The  presence  of 
FeCl3  or  heating  the  solution  does  not  reduce  the  As2S5  to  As2S3  .  If  there 
be  a  small  amount  of  HC1  and  the  H2S  be  passed  in  slowly  about  15  per 
cent  of  As2S3  is  formed:  2H3As04  +  5H2S  +  xHCl  =  As2S3  +  S2  -f 
8H20  +  xHCl  .  If  NH4C1  be  present  more  As2S3  is  formed.  According 
to  Thiele  (C.  C.,  1890,  1,  877),  arsenic  acid  cold  treated  with  a  slow  stream 
of  H2S  gives  arsenous  sulphide,  while  the  hot  acid  with  a  rapid  stream  of 
the  gas  gives  the  pentasulphide.  Arsenic  sulphide  has  the  same  solubili- 
ties as  arsenous  sulphide.  When  distilled  with  hydrochloric  acid  gas 
and  a  reducing  agent,  arsenous  chloride  is  formed  (AsCl5  is  not  known 
to  exist).  The  solutions  in  the  alkali  hydroxides,  carbonates  and  sul- 
phides form  arsenates  and  thioarsenates  (5c).  Ammonium  sulphide  added 
to  a  neutral  or  alkaline  solution  of  arsenic  acid  forms  arsenic  sulphide 
which  remains  in  solution  as  ammonium  thioarsenate  (5c).  The  addition 
of  acid  at  once  forms  arsenic  sulphide,  not  arsenous  sulphide  and  sulphur. 
The  reaction  is  much  more  rapid  than  with  hydrosulplmric  acid  and  is 
facilitated  by  warming. 

Arsine,  AsH  .  ,  does  not  combine  with  hydrosulphuric  acid  until  heated 
to  230°,  while  stibine,  SbH3  ,  combines  at  the  ordinary  temperature  (Brunn, 
B.,  1889,  22,  3202). 

Acidulated  solutions  of  arsenic  boiled  with  thiosulphates  form  arsenous 
sulphide  (separation  f  rom  Sb  and  Sn)  (Lesser,  Z.,  1888,  27,  218).  Arsenic 
may  be  removed  from  sulphuric  acid  by  boiling  with  barium  thiosulphate 
and  no  foreign  material  is  introduced  into  the  acid:  As203  -f-  3BaS,0:,  = 
As2S3  +  3BaS04  ;  2H3As04  +  5Na2S203  =  As2S3  +  5Na2S04  +  S2  +  3H20. 
(Thorn,  J.  C.,  1876,  29,  517;  Wagner,  Dingl,  1875,  218,  321). 

Sulphurous  acid  readily  reduces  arsenic  acid  to  arsenous  acid  :  H3As04  + 
H2S03  =  H3As03  +  H2S04  (Woehler,  A.,  1839,  30,  224). 


§69,  Qi.  ARSENIC-  61 

f. — The  arsenic  from  all  arsenical  compounds  treated  with  concentrated 
hydrochloric  acid  and  then  distilled  in  a  current  of  hydrochloric  acid  gas, 
passes  into  the  distillate  as  arsenous  chloride,  AsCl3  .  Nearly  all  of  the 
arsenic  will  be  carried"  over  in  the  first  50  cc.  of  the  distillate.  This  is  a 
very  accurate  quantitative  separation  of  arsenic  from  antimony  and  tin 
and  from  other  non-volatile  organic  and  inorganic  material.  The  AsCl3 
passes  over  at  132°,  condenses  with  HC1  and  may  be  tested  with  SnCl2 
(g),  or,  after  decomposition  with  water  (5c)  by  the  usual  tests  for  arsenous 
acid  (Huf schmidt,  B.,  1884,  17,  2245;  Beckurts,  Arch.  Pharm.,  1884,  222, 
684;  Piloty  and  Stock,  B.,  1897,  30,  1649). 

Hydrobromic  acid  in  dilute  solutions  is  without  action  upon  the  acids 
of  arsenic.  The  concentrated  acid  reduces  arsenic  acid  to  arsenous  acid: 
H3As04  +  2HBr  -  ~  H,As03  -f  Br,  +  H20  .  Hydriodic  acid  reduces 
arsenic  acid  to  arsenous  acid  with  liberation  of  iodine.  This  is  a  method 
of  detecting  Asv  in  the  presence  of  As'".  0.0001  gram  of  H3As04  may  be 
detected  in  the  presence  of  one  gram  of  As203  :  2H3As04  +  4HI  =  As203 
+  2I2  +  5H20  (Naylor,  J.  C.,  1880,  38,  421). 

Chloric  and  bromic  acids  oxidize  arsenous  compounds  to  arsenic  acid  with 
formation  of  the  corresponding  hydracid:  3As203  +  2HBr03  -f~  9H20  = 
6HS  AsO4  +  2HBr .  lodic  acid  oxidizes  arsenous  compounds  to  arsenic  acid 
with  liberation  of  iodine:  5As203  +  4HI03  +  13H20  =  10H3As04  +  2I2  . 

g. — Stannous  chloride,  SnCL  ,  reduces  all  compounds  of  arsenic  from  their 
hot  concentrated  hydrochloric  acid  solutions,  as  flocculent,  black-brown,  metal- 
loidal  arsenic,  containing  three  or  four  per  cent  of  tin.  The  arsenic,  in  solution 
with  the  concentrated  hydrochloric  acid,  acts  as  arsenous  chloride:  4AsCl3  -f- 
GSnCl,  —  As4  +  6SnCl4  .'  The  hydrochloric  acid  should  be  25  to  33  per  cent;  if 
not  over  15  to  20  per  cent,  the  reaction  is  slow  and  imperfect. 

In  a  wide  test-tube  place  0.1  to  0.2  gram  of  the  (oxidized)  solid  or  solution 
to  be  tested,  add  about  1  gram  of  sodium  chloride,  and  2  or  3  cc.  of  sulphuric 
acid,  then  about  1  gram  of  crystallized  stannous  chloride',  agitate,  and  heat  to 
boiling  several  times,  and  set  aside  for  a  few  minutes.  Traces  of  arsenic  give 
only  a  brown  color;  notable  proportions  give  the  flocculent  precipitate.  A 
dark  gray  precipitate  may  be  due  to  mercury  (§58,  60),  capable  of  being  gath- 
ered into  globules.  If  a  precipitate  or  a  darkening  occurs,  obtain  conclusive 
evidence  whether  it  contains  arsenic  or  not,  as  follows:  Dilute  the  mixture 
with  ten  to  fifteen  volumes  of  about  12  per  cent  hydrochloric  acid;  set  aside, 
decant;  gather  the  precipitate  in  a  wet  filter,  wash  it  with  a  mixture  of  hydro- 
chloric acid  and  alcohol,  then  with  alcohol,  then  with  a  little  ether,  and  dry  in 
a  warm  place.  A  portion  of  this  dry  precipitate  is  now  dropped  into  a  small 
hard-glass  tube,  drawn  out  and  closed  at  one  end,  and  heated  in  the  flame; 
arsenic  is  identified  by  its  mirror  (7),  easily  distinguished  from  mercury 
(§58,  7).  Antimony  is  not  reduced  by  stannous  chloride;  other  reducible 
metals  give  no  mirror  in  the  reduction-tube.  Small  proportions  of  organic 
material  impair  the  delicacy  of  this  reaction,  but  do  not  prevent  it.  It  is 
especially  applicable  to  the  hydrochloric  acid  distillate,  obtained  in  separation 
of  arsenic,  according  to  f. 

h. — Chromates  boiled  with  arsenites  and  sodium  bicarbonate  give  chromium 
arsenatc  (Tarugi,  J.  (7.,  1896,  70,  ii,  340  and  390). 

i. — Magnesium  salts  with  ammonium  chloride  and  ammonium  hydroxide 
precipitate  from  solutions  of  arsenates,  magnesium  ammonium  arsenate, 
MgNH,AsO4  ,  white,  easily  soluble  in  acids.  The  reagents  should  be  first 
mixed  together,  and  used  in  a  clear  solution  ("  magnesia  mixture  ")  to  make 
sure  that  enough  ammonium  salt  is  present  to  prevent  the  precipitation  of 
magnesium  hydroxide  by  the  ammonium  hydroxide.  The  crystalline  precipi- 


62  ARSENIC.  §69,  6;. 

tate  forms  slowly  but  completely.  Compare  with  the  corresponding  magnesium 
ammonium  phosphate  (§189,  Qd).  Maync-xium  urxcnite  is  insoluble  in  water,  but 
is  soluble  in  ammonium  hydroxide  and  in  ammonium  chloride  (distinction  from 
arsenates). 

j. — Silver  nitrate  solution  precipitates  from  neutral  solutions  of  arsenites,  or 
ammonio-silver  nitrate  *  precipitates  from  a  water  solution  of  arsenous  oxide, 
silver  arsenite,  Ag3AsO"3  ,  yellow,  readily  soluble  in  dilute  acids  or  in  ammonium 
hydroxide  (§59,  6#).  Neutral  solutions  of  arsf.  nates  are  precipitated  as  silver 
arsenate,  Ag3As04  ,  reddish  brown,  having  the  same  solubilities  as  the  arsenite. 

k. — Copper  sulphate  solution  precipitates  from  neutral  solutions  of  arsenites, 
or  ammonio-copper  sulphate  (prepared  in  the  same  manner  as  the  ammonio- 
silver  oxide  described  above)  precipitates  from  water  solutions  of  arsenous 
oxide,  the  ffreen  copper  arsenitc,  CuHAsO3  (Scheele's  green),  soluble  in  ammo- 
nium hydroxide  and  in  dilute  acids.  Copper  acetate,  in  boiling  solution,  pre- 
cipitates the  green  copper  aceto-arsenite  (CuOAs2O3)3Cu(C2H3Oa)2  (Scliweinfurt 
green),  soluble  in  ammonium  hydroxide  and  in  acids.  Both  these  salts  are 
often  designated  as  Paris  green  (§77,  6#).  Copper  sulphate  with  excess  of  free 
alkali  is  reduced  to  cuprous  oxide  with  formation  of  alkali  arsenate  (10). 
K3As03  +  2CuS04  +  4KOH  =  K3As04  +  2K2S04  +  Cu2O  +  2H2O  .  Solutions 
of  arsenates  are  precipitated  by  copper  sulphate  as  copper  arsenate,  CuHAsO,  , 
greenish  blue,  the  solubilities  and  conditions  of  precipitation  being  the  same 
as  for  the  arsenites. 

1. — Ferric  salts  precipitate  from  arsenites,  and  freshly  precipitated  ferric 
hydroxide  (used  as  an  antidote,  Wormley,  246),  forms  with  arsenous  oxide, 
variable  basic  ferric  arsenites,  scarcely  soluble  in  acetic  acid,  soluble  in  hydro- 
chloric acid.  Water  slowly  and  sparingly  dissolves  from  the  precipitate  the 
arsenous  anhydride;  but  a  large  excess  of  the  ferric  hydroxide  holds  nearly  all 
the  arsenic  insoluble.  To  some  extent  the  basic  ferric  arsenites  are  trans- 
posed into  basic  ferrous  arsenates,  insoluble  in  water,  in  accordance  with  the 
reducing  power  of  arsenous  oxide.  In  the  presence  of  alkali  acetates,  arsenic 
acid,  or  acidulated  solutions  of  arsenates,  are  precipitated  by  ferric  salts  as  ferric 
arsenate,  FeAsO4,  yellowish  white,  insoluble  in  acetic  acid  (compare  §126,  6ef). 

m. — Ammonium  molybdate,  (NH4)2MoO4,  in  nitric  acid  solution,  when  slightly 
warmed  with  a  solution  of  arsenic  acid  or  of  arsenates  gives  a  yellow  precipitate  of 
ammonium  arseno-molybdale,  of  variable  composition.  No  precipitate  is  formed  with 
As'".  This  precipate  is  very  similar  in  appearance  and  properties  to  the  ammonium 
phospho-molybdate;  except  the  latter  precipitates  completely  in  the  cold. 

6'.  Special  Reactions,  a. — Marsh's  Test. — This  is  an  extremely  deli- 
cate test  for  arsenic,  especially  adapted  for  the  detection  of  this  ele- 
ment when  present  in  small  quantities,  even  when  large  quantities  of 
other  elements  are  present.  Arsenic,  from  all  of  its  soluble  com- 
pounds, is  reduced  by  the  action  of  dilute  sulphuric  or  hydrochloric 
acid  on  zinc,  forming  at  first  metallic  arsenic  and  then  arsenous  hydride, 
AsH,,  gaseous:  As203  +  6Zn  +  6H2S04  =  2AsH3  +  6ZnS04  +  3H20  ; 
H3As04  +  4Zn  +  4H2S04  =±  AsH3  +  4ZnS04  +  4H20  .  The  arsenic  is 
precipitated  with  the  other  metals  of  the  second  group  by  hydrogen 
sulphide,  separated  with  antimony,  tin  (gold,  platinum  and  molybdenum) 
by  yellow  ammonium  sulphide.  This  solution  is  precipitated  by  dilute 
hydrochloric  acid  and  the  mixed  sulphides,  well  washed^  are  dissolved  in 
hydrochloric  acid  using  as  small  an  amount  of  potassium  chlorate  crystals 
as  possible.  The  solution  is  boiled  (till  it  does  not  bleach .  litmus  paper) 

*  Prepared  by  adding  ammonium  hydroxide  to  a  solution  of  silver  nitrate  till  the  precipitate 
at  first  produced  is  marly  all  redissoived. 

T  If  the  ammonium  salts  are  not  thoroughly  removed  by  washing  there  is  danger  of  the  for- 
mation of  the  very  explosive  chloride  of  nitrogen  (§268,  1)  when  the  precipitate  is  treated 
with  hydrochloric  acid  and  potassium  chlorate. 


CALIFORWW   COUEfil 

u**ta.      °*  P"ARMACY   63 


to  remove  excess  of  chlorine  and  is  then  ready  for  the  Marsh  apparatus. 
This  apparatus  consists  of  a  strong  Erlenmeyer  flask  of  about  125  cc. 
capacity  fitted  with  a  two  hole  rubber  stopper.  Through  one  hole  is  passed 
a  thistle  (safety)  tube,  reaching  nearly  to  the  bottom  of  the  flask;  in  the 
other  is  fitted  a  three-inch  Marchand  calcium  chloride  tube,  which  projects 
just  through  the  stopper  and  is  filled  with  glass-wool  and  granular  calcium 
chloride  to  dry  the  gases  generated  in  the  flask.  To  the  other  end  of 
the  Marchand  tube  is  fitted,  with  a  small  cork  or  rubber  stopper,  a  piece 
of  hard  glass  tubing  of  six  mm.  diameter  and  one  foot  long.  This  tube 
should  be  constricted  one-half,  for  about  two  inches,  beginning  at  the 
middle  of  the  tube  and  extending  toward  the  end  not  fastened  to  the 
calcium  chloride  tube.  The  outer  end  of  the  tube  should  also  be  con- 
stricted to  about  one  mm.  inner  diameter.  A  short  piece  of  rubber  tubing 
should  connect  this  constricted  end  with  a  piece  of  ordinary  glass  tubing, 
dipping  into  a  test  tube  about  two-thirds  filled  with  a  two  per  cent  solu- 
tion of  silver  nitrate.  The  rubber  tubing  should  make  a  close  joint  with 
the  constricted  end  of  the  hard  glass  tube,  and  yet  not  fit  so  snug  but  that 
it  can  be  easily  removed. 

From  10  to  20  grams  of  granulated  zinc  *  are  placed  in  the  flask  with 
sufficient  water  to  cover  the  end  of  the  thistle  tube.  Four  or  five  cubic 
centimeters  of  reagent  sodium  carbonate  are  added  and  the  stopper 
tightly  fitted  to  the  flask.  Dilute  sulphuric  acid  (one  of  acid  to  three  of 
water)  should  now  be  added,  very  carefully  at  first,  f  until  a  moderate 
evolution  of  hydrogen  is  obtained,. 

The  hydrogen  should  be  allowed  to  bubble  through  the  silver  nitrate 
for  about  five  minutes.  There  should  be  no  appreciable  blackening  of 
the  solution  (§59,  10),  thus  proving  the  absence  of  arsenic  from  the  zinc 
and  the  sulphuric  acid.  The  purity  of  the  reagents  having  been  estab- 
lished the  solution  containing  the  arsenic  may  be  added  in  small  amounts 
at  a  time  through  the  thistle  tube.  If  arsenic  be  present  there  will  be 
almost  immediate  blackening  of  the  silver  nitrate  solution. 

6AgN03  +  AsH3  +  3H,0  =  6Ag  +  H3As03  +  GHN03 

The  hard  glass  tube  should  now  be  heated  J  to  redness  by  a  flame  from 

*  The  zinc  and  all  the  reagents  should  be  absolutely  free  from  arsenic.  If  the  zinc  be  strictly 
chemically  pure  it  will  be  but  slowly  attacked  by  the  acid.  It  should  be  platinized  (§349,  4a)or 
should  contain  traces  of  iron.  Hote  '  A.  Ch.,  1884,  (6>,  3,  141)  removes  arsenic  from  zinc  by  adding 
anhydrous  MgCl4  to  the  molten  metal,  A»C13  being  evolved.  The  zinc  purified  in  this  way  is 
readily  attacked  by  acids. 

t  The  acid  first  added  decomposes  the  alkali  carbonate  forming  carbon  dioxide  which  rapidly 
displaces  the  air  and  greatly  lessens  the  danger  of  explosion  when  the  gas  is  ignited.  If  too 
much  acid  be  added  before  the  carbonate  is  decomposed  violent  frothing  may  take  place  and 
the  liquid  contents  of  the  flask  forced  into  the  calcium  chloride  tube. 

\  Before  heating  the  tube  or  igniting  the  gas,  a  towel  should  be  wrapped  around  the  flask  to 
insure  safety  in  case  of  an  explosion  due  to  the  imperfect  removal  of  the  air  ;  or  the  tube  con- 
necting the  hard  glass  tube  with  the  Marchand  tube  should  be  of  larger  size  and  provided  with 
u  plug  of  wire  gauze  (made  of  10  or  20  circles  of  gauze  the  size  of  the  tube).  A  flame  cannot 
pass  such  a  plug  of  wire  gauze. 


64  AlttiKXic.  §69, 6'fe. 

a  Bunsen  burner  provided  with  a  flame  spreader.  The  flame  should  be 
applied  to  the  tube  between  the  calcium  chloride  tube  and  the  constricted 
portion.  The  tube  should  be  supported  to  prevent  sagging  in  case  the 
glass  softens,  and  it  is  customary  to  wrap  a  few  turns  of  wire  gauze  around 
the  portion  of  the  tube  receiving  the  heat.  The  heat  of  the  flame  decom- 
poses the  arsine  and  a  mirror  of  metallic  arsenic  is  deposited  in  the  con- 
stricted portion  of  the  tube  just  beyond  the  heated  portion.  This  may 
be  tested  as  described  under  c  1.  When  a  sufficient  mirror  has  been 
obtained  the  flame  is  withdrawn,  and,  removing  the  rubber  tube,  the 
escaping  gas  *  is  ignited.  As  small  a  quantity  of  arsenic  as  0.002  mg. 
will  produce  a  visible  mirror  and  if  20  g.  of  material  is  used  for  analysis, 
this  would  represent  one  part  of  arsenic  in  10,000,000  parts. 

1).  Arsenous  Hydride  (arsine),  AsH3 ,  burns  when  a  stream  of  it  is  ignited 
where  it  enters  the  air,  and  explodes  when  its  mixture  with  air  is  ignited. 
It  burns  with  a  somewhat  luminous  and  slightly  bluish  flame  (distinction 
from  hydrogen)  ;  the  hydrogen  being  first  oxidized,  and  the  liberated 
arsenic  becoming  incandescent,  and  then  undergoing  oxidation;  the  vapors 
of  water  and  arsenous  anhydride  passing  into  the  air:  2AsH3  +  302  = 
As203  -f  3H20  .  If  present  in  considerable  quantity  a  white  powder  may 
be  observed  settling  on  a  piece  of  black  paper  placed  beneath  the  flame. 
If  the  cold  surface  of  a  porcelain  dish  be  brought  in  contact  with  the 
flame  the  oxidation  is  prevented  and  lustrous  black  or  brownish-black 
spots  of  metallic  arsenic  are  deposited  on  the  porcelain  surface;  4AsH3  + 
302  =  As4  +  GH.,0  .  A  number  of  spots  should  be  obtained  and  all  the 
tests  for  metallic  arsenic  applied.  The  arsenic  in  the  silver  nitrate  solu- 
tion is  present  as  arsenous  acid  and  can  be  detected  by  the  usual  tests  (6e) 
by  first  removing  the  excess  of  silver  nitrate  with  dilute  hydrochloric  acid 
or  calcium  chloride,  or  by  cautiously  neutralizing  with  ammonia  the 
arsenic  may  be  precipitated  as  the  yellow  silvery  arsenite  (6/). 

To  generate  arsine,  magnesium  or  iron  f  may  be  used,  instead  of  zinc,  and 
hydrochloric  acid  instead  of  sulphuric  acid.  Arsine  cannot  be  formed  in  the 
presence  of  oxidizing  agents  as  the  halogens,  nitric  acid,  chlorates,  hypo- 
chlorites,  etc.  Arsinuretted  hydrogen  (arsine)  may  also  be  produced  from 
arsenous  compounds  by  nascent  hydrogen  generated  in  alkaline  solution.  Sodium 
amalgam,  t  zinc  (or  zinc  and  magnesium)  and  potassium  hydroxide  or  alumi- 
num and  potassium  hydroxide  may  be  used  as  the  reducing  agent.  There  is 

*  Arsine  is  an  exceedingly  poisonous  gas,  the  inhalation  of  the  unmixed  gas  being  quickly 
fatal.  Its  dissemination  in  the  air  of  the  laboratory,  even  in  the  small  portions  which  are  not 
appreciably  poisonous,  should  be  avoided.  Furthermore,  as  it  is  recognized  or  determined,  in 
its  various  analytical  reactions,  only  by  its  decomposition,  to  permit  it  to  escape  undecomposed 
is  so  far  to  fail  in  the  object  of  its  production.  The  evolved  gas  should  be  constantly  run  into 
silver  nitrate  solution,  or  kept  burning. 

t  According  to  Thiele  (C.  C.,  1890, 1,  877)  arsenic  may  be  separated  from  antimony  in  the  Marsh 
test  by  using  electrolytirally  deposited  iron  instead  of  zinc.  Stibine  is  not  evolved.  According 
to  Sautermeister  (Analyst,  1891,  218)  arsine  is  not  produced  when  hydrochloric  acid  acts  upon 
iron  containing  arsenic,  but  if  several  grams  of  zinc  be  added  a  very  small  amount  of  arsenic  in 
the  iron  may  be  detected. 

%  Sodium  amalgam  is  conveniently  prepared  by  adding  (in  small  pieces  at  a  tiroo)  one  part  of 
sodium  to  eight  parts  (by  weight)  of  dry  mercury  warmed  on  the  water  bath.  When  cold  the 
amalgam  becomes  solid  and  is  easily  broken.  It  should  be  preserved  in  well  stoppered  bottles. 


§69, 6'c.  ARSENIC.  6*5 

no  reaction  with  Asv,  or  with  compounds  of  antimony  (§70,  6;');  hence  when 
the  arsenic  is  present  in  the  triad  condition  (Asv  may  be  reduced  to  As"'  by 
SO.,)  the  use  of  one  of  the  above  reagents  serves  aclinirably  i'or  the  detection 
of  arsenic  in  the  presence  of  antimony.  This  experiment  may  be  made  in  a 
test-tube,  the  arsenic  being  detected  by  covering  the  tube  with  a  piece  of  filter 
paper  moistened  with  silver  nitrate.  It  is  very  difficult  to  drive  over  the  last 
traces  of  the  arsenic  and  therefore  the  method  is  not  satisfactory  for  quanti- 
tative work  (Hager,  /.  C.,  1885,  48,  838;  Johnson,  C.  N.,  1878,  38,  301;  and  Clark, 
J.  C.,  1893,  63,  884). 

If  ferrous  sulphide  contains  metallic  iron  and  arsenic,  arsine  may  be  gen- 
erated with  the  hydrogen  sulphide.  It  cannot  be  removed  by  washing  the 
gases  with  hydrochloric  acid  (Otto,  B.,  1883,  16,  2947). 

Arsine  does  not  combine  with  hydrogen  sulphide  until  heated  to  230°,  while 
.v/M/w,  SbH3  ,  combines  at  ordinary  temperature  (method  of  separation) 
(Hrunn,  B.,  1889,  22,  3202;  Myers,  /.  C.,  1871,  24,  889).  As  dry  hydrogen  sul- 
phide is  without  action  upon  dry  iodine,  it  may  be  freed  from  arsine  by  passing 
the  mixture  of  the  dried  gases  through  a  tube  filled  with  glass  wool  inter- 
spersed with  dry  iodine.  AsHs  +  3la  =  Asl,  -f  SHI  (Jacobson,  B.,  1887,  20, 
1999).  Arsenous  hydride  is  decomposed  by  passing  through  a  tube  heated  to 
redness  (mirror  in  Marsh  test)  4AsHg  =  As4  +  6H2  .  Nitric  acid  oxidizes  it 
to  arsenic  acid,  3AsH,  -f  8HNO3  =  3H8AsO4  -f  8NO  +  4H,O;  and  may  be  used 
instead  of  silver  nitrate  to  effect  a  separation  of  arsine  and  stibine  in  the 
Marsh  test.  The  nitric  acid  solution  is  evaporated  to  dryness  and  the  residue 
thoroughly  washed  with  water.  Test  the  solution  for  arsenic  with  silver 
nitrate  and  ammonium  hydroxide  (Ag8AsO4  ,  reddish  brown  precipitate,  &j). 
Dissolve  the  residue  in  hydrochloric  or  nitrohydrochloric  acid  and  test  for 
antimony  with  hydrogen  sulphide  (Ansell,  J.  0.,  1853,  5,  210). 

c. — Comparison  of  the  mirrors  and  spots  obtained  with  arsenic  and  anti- 
mony.— 1.  Both  the  mirror  and  spots  obtained  in  the  Marsh  test  exhibit 
the  properties  of  elemental  arsenic  (5a).  The  reactions  of  these  deposits 
having  analytical  interest  are  such  as  distinguish  arsenic  from  antimony. 

ARSENIC  MIRROR.  ANTIMONY  MIRROR. 

Deposited  beyond  the  flame;  ar-  Deposited  before  or  on  both  sides 
sene  not  being  decomposed  much  be-  of  the  flame;  stibine  being  decom- 
low  a  red  heat.  poied  considerably  below  a  red  heat. 

Volatilizes  in  absence  of  air  at  The  mirror  melts  to  minute  glob- 

450°  (1),  allowing  the  mirror  to  be  ules  at  630°,  and  is  then  driven  at 

driven  along  the  tube;  it  does  not  a  red  heat, 
melt. 

By  vaporization  in  the  stream  of          The  vapor  has  no  odor, 
gas,  escapes  with  a  garlic  odor. 

By  slow  vaporization  in  a  cur-  By  vaporization  in  a  current  of 
rent  of  air  a  deposit  of  octahedral  air,  a  .white  amorphous  coating  is 
and  tetrahedral  crystals  is  obtained,  obtained;  insoluble  in  water,  soluble 
forming  a  white  coating  soluble  in  in  hydrochloric  acid,  and  giving  re- 
water  and  giving  the  reactions  for  actions  for  antimonous  oxide, 
arsenous  oxide. 


GG 


ARSENIC. 


§69,  6><?> 


The  heated  mirror  combines  with 
hydrogen  sulphide,  forming  the 
lemon-yellow  arsenous  sulphide, 
which,  being  volatile,  is  driven  to 
the  cooler  portion  of  the  tube. 

The  dry  sulphide  is  not  readily 
attacked  by  dry  hydrochloric  acid 
gas  (6/). 


Arsenic  Spots. 
Of  a  steel  gray  to  black  lustre. 

Volatile  by  oxidation  to  arsenous 
oxide  at  218°. 

Dissolve  in  hypochlorite.* 

Warmed  with  a  drop  of  ammon- 
ium sulphide  form  yellow  spots, 
soluble  in  ammonium  carbonate,  in- 
soluble in  hydrochloric  acid  (6e). 


With  a  drop  of  hot  nitric  acid, 
dissolve  clear.  The  clear  solution, 
with  a  drop  of  solution  of  silver 
nitrate,  when  treated  with  vapor  of 
ammonia,  gives  a  brick-red  precipi- 
tate. 


The  solution  gives  a  yellow  pre- 
cipitate when  warmed  with  a  drop 
of  ammonium  molybdate. 

With  vapor  of  iodine,  color  yel- 
low, by  formation  of  arsenous 
iodide,  readily  volatile  when  heated. 


The  heated  mirror  combines  with 
hydrogen  sulphide  forming  the 
orange  antimonous  sulphide,  which 
is  not  readily  volatile. 


The  sulphide  is  readily  decom- 
posed by  dry  hydrochloric  acid  gas, 
forming  antimonous  chloride  which 
is  volatile,  and  may  be  driven  over 
the  unattacked  arsenous  sulphide. 


Antimony 
Of  a  velvety  brown  to  black  sur- 


face. 

Volatile,    by    oxidation    to 
monous  oxide,  at  a  red  heat. 


anti- 


Do  not  dissolve  in  hypochlorite. 

Warmed  with  ammonium  sul- 
phide, form  orange-yellow  spots,  in- 
soluble in  ammonium  carbonate, 
soluble  in  hydrochloric  acid  (§70. 
Qe). 

With  a  drop  of  hot  dilute  nitric 
acid,  turn  white.  The  white  fleck. 
by  action  of  nitric  acid  treated  with 
silver  nitrate  and  vapor  of  ammo- 
nia, gives  no  color  until  warmed 
with  a  drop  of  ammonium  hydrox- 
ide, then  gives  a  black  precipitate. 

With  the  white  fleck  no  further 
action  on  addition  of  ammonium 
molybdate. 

With  vapor  of  iodine,  color  more 
or  less  carmine-red,  by  formation 
of  antimonous  iodide,  not  readily 
volatile  by  heat. 


*The  hypochlorite  reagent,  usually  NaCIO,  decomposes  in  the  air  and  light  on  standing. 
It  should  instantly  and  perfectly  bleach  litmus  paper  (not  redden  it).  It  dissolves  arsenic  by 
oxidation  to  arsenic  acid.  As4  -f  lOXaCIO  +  6H2O  =  4H3A«O4  +  lONaCl. 


£69, 6'd.  ARSENIC.  67 

2.  To  the  spot  obtained  on  the  porcelain  surface,  add  a  few  drops  of 
nitric  acid  and  heat;  then  add  a  drop  of  ammonium  molybdate.     A  yellow 
precipitate   indicates   arsenic.     Antimony  may  give   a  white  precipitate 
with  the  nitric  acid,  but  gives  no  further  change  with  the  ammonium 
molybdate  (Deniges,  C.  r.,  1890,  111,  824). 

3.  Oxidize  the  arsenic  spot  with  nitric  acid  and  evaporate  to  dryness. 
Add  a  drop  of  silver  nitrate  or  ammonio-silver  nitrate  (6/).     A  reddish- 
brown  precipitate  indicates  arsenic. 

4.  After  the  formation  of  the  mirror  in  Marsh's  test  the  generating 
flask  may  be  disconnected  and  a  stream  of  dry  hydrogen  sulphide  passed 
over  the  heated  mirror.     If  the  mirror  consists  of  both  arsenic  and  anti- 
mony, the  sulphides  of  both  these  metals  will  be  formed,   and  as  the 
arsenous  sulphide  is  volatile  when  heated,  it  will  be  deposited  in  the  cooler 
portion  of  the  tube.     The  sulphides  being  thus  separated  can  readily  be 
distinguished  by  the  color.     If  now  a  current  of  dry  hydrochloric  acid 
gas  be  substituted  for  the  hydrogen  sulphide  the  antimonous  sulphide 
will  be  decomposed  to  the  white  antimonous  chloride  which  volatilizes  and 
may  be  driven  past  the  unchanged  arsenous  sulphide  (5c). 

5.  The  tube  containing  the  mirror  is  cut  so  as  to  leave  about  two  inches 
on  each  side  of  the  mirror  and  left  open  at  both  ends.     Incline  the  tube 
and  beginning  at  the  lower  edge  of  the  mirror  gently  heat,  driving  the 
mirror  along  the  tube.     The  mirror  will  disappear  and  if  much  arsenic 
be  present  a  white  powder  will  be  seen  forming  a  ring  just  above  the 
heated  portion  of  the  tube.     This  powder  consists  of  crystals  of  arsenous 
oxide,  and  should  be  carefully  examined  under  the  microscope  and  iden- 
tified by  their  crystalline  form  (Wormley,  270). 

6.  The  crystals  of  arsenous  oxide  obtained  above  are  dissolved  in  water 
and  treated  with  ammonio-silver  nitrate  forming  the  yellow  silver  arse- 
nil  c  (C)/):  or  with   ammonio-copper  sulphate  forming  the  green  copper 
ai-smite  (6fc)  (Wormley,  259).     Any  other  test  for  arsenous  oxide  may  be 
applied  as  desired. 

7.  Magnesia  mixture  (Gi)  is  added  to  the  solution  of  the  mirror  or  spots 
in  nitric  acid.    The  solution  must  be  strongly  alkaline.    A  white  crystalline 
precipitate   of  magnesium  ammonium  arsenate,   MgNH4As04,   is   formed 
(Wormley,  316). 

d. — Remsch's  Test. — If  a  solution  of  arsenic  be  boiled  with  hydrochloric  acid 
and  a  strip  of  bright  copper  foil,  the  arsenic  is  deposited  on  the  copper  as  a 
gray  film.  Hager  (C.  C.,  1886,  680)  recommends  the  use  of  brass  foil  instead  of 
copper  foil.  When  a  large  amount  of  arsenic  is  present  the  coating  of  arsenic 
separates  from  the  copper  in  scales.  The  film  does  not  consist  of  pure  metallic 
arsenic,  but  appears  to  be  an  alloy  of  arsenic  and  copper.  Arsenous  compounds 
are  reduced  much  more  readily  than  arsenic  compounds.  The  hydrochloric 
acid  should  compose  at  least  one-tenth  the  volume  of  the  solution.  The  arsenic 
is  not  deposited  if  the  acid  is  no*  present.  This  serves  as  one  of  the  most 
satisfactory  methods  of  determining  the  presence  or  absence  of  arsenic  jn 


68  AR8ENIC.  §69,  6'g. 

hydrochloric  acid.  Dilute  the  concentrated  acid  with  five  parts  of  water  and 
boil  with  a  thin  strip  of  bright  copper  foil.  A  trace  of  arsenic  if  present  will 
soon  appear  on  the  foil.  For  further  identification  of  the  deposit,  wash  the 
foil  with  distilled  water,  dry,  and  heat  in  a  hard  glass  tube,  as  for  the  oxida- 
tion of  the  arsenic  mirror  (6'c,  5).  The  crystals  may  be  identified  by  the  mic- 
roscope and  by  any  other  tests  for  arsenous  oxide.  It  is  important  that  the 
surface  of  the  copper  should  be  bright.  This  is  obtained  by  rubbing  the  sur- 
face of  the  foil  with  a  file,  a  piece  of  pumice  or  sand-paper  just  before  using. 
The  copper  should  not  contain  arsenic,  but  if  it  does  contain  a  small  amount 
no  film  will  be  deposited  due  to  its  presence  unless  agents  are  present  which 
cause  partial  solution  of  the  foil.  If  a  strip  of  the  foil,  upon  boiling  with 
hydrochloric  acid  for  ten  minutes,  shows  no  dimming  of  the  brightness  of 
the  copper  surface;  the  purity  of  both  acid  and  copper  may  be  relied  upon  for 
the  most  exact  work.  Antimony,  mercury,  silver,  bismuth,  platinum,  palladium 
and  gold  are  deposited  upon  copper  when  boiled  with  hydrochloric  acid.  Under 
certain  conditions  most  of  these  deposits  ma}r  closely  resemble  that  of  arsenic. 
Of  these  metals  mercury  is  the  only  one  that  forms  a  sublimate  when  heated 
in  the  reduction  tube  (7),  and  this  is  readily  distinguished  from  arsenic  by 
examination  under  the  microscope.  Antimony  may  be  volatilized  as  an  amor- 
phous powder  at  a  very  high  heat.  Organic  material  may  sometimes  give  a 
deposit  on  the  copper  wrhich  also  yields  a  sublimate,  but  this  is  amorphous  and 
does  not  show  the  octahedral  crystals  wrhen  examined  under  the  microscope 
(Wormley,  269  and  ff.;  Clark,  J.  (7.,  1893,  63,  886). 

€.• — Detection  in  Case  of  Poisoning1. — Arsenic  in  its  various  compounds  is 
largely  used  as  a  poison  for  bugs,  rodents,  etc.,  and  frequently  cases  arise  of 
accidental  arsenical  poisoning.  It  is  also  used  for  intentional  poisoning,  chiefly 
suicidal.  It  is  usually  taken  in  the  form  of  arsenous  oxide  (white  arsenic),  or 
"  Fowler's  Solution  "  (a  solution  of  the  oxide  in  alkali  carbonate).  One  hun- 
dred fifty  to  two  hundred  milligrams  (two  to  three  grains)  are  usually  sufficient 
to  produce  death.  Violent  vomiting  is  a  usual  symptom  and  death  occurs  in 
from  three  to  six  hours.  In  cases  of  suspected  poisoning  vomiting  should  be 
induced  as  soon  as  possible  by  using  an  emetic  followed  by  demulcent  drinks, 
or  the  stomach  should  be  emptied  by  a  stomach  pump.  Freshly  prepared  ferric 
hydroxide  is  the  usual  antidote,  of  which  twenty-five  to  fifty  grams  (one  to 
two  ounces)  may  be  given.  The  antidote  may  be  prepared  by  adding  magnesia 
(magnesium  oxide),  ammonium  hydroxide,  or  cooking  soda  (sodium  bicarbo- 
nate) to  ferric  chloride  or  muriate  tincture  of  iron:  straining  in  a  clean  piece 
of  muslin,  and  washing  several  times.  If  magnesia  be  used  it  is  not  necessary 
to  wash,  as  the  magnesium  chloride  formed  is  helpful  rather  than  injurious. 
A  portion  of  the  ferric  hydroxide  oxidizes  some  of  the  arsenous  compound, 
being  itself  reduced  to  the  ferrous  condition,  and  forming  an  insoluble  ferrous 
arsenate.  When  the  ferric  oxide  is  in  excess  the  ferrous  arsenate  does  not 
appear  to  be  acted  upon  by  the  acids  of  the  stomach.  Of  course  it  will  be  seen 
that  the  ferric  hydroxide  will  have  no  effect  upon  the  arsenic  which  has 
entered  into  the  circulation. 

It  frequently  becomes  necessary  for  the  chemist  to  analyze  portions  of  sus- 
pected food,  contents  of  the  stomach,  urine;  or,  if  death  has  ensued,  portions 
of  the  stomach,  intestines,  liver,  or  other  parts  of  the  body.  At  first  a  careful 
examination  should  be  made  of  the  material  at  hand  for  solid  white  particles, 
that  would  indicate  arsenous  oxide.  If  particles  be  found  they  can  at  once  be 
identified  by  the  usual  tests.  Liquid  food  or  liquid  contents  of  the  stomach 
should  be  trailed  with  dilute  hydrochloric  acid,  filtered  and  washed  and  the 
filtrate  precipitated  with  hydrogen  sulphide,  etc.  When  solid  food  or  portions 
of  tissue  are  to  be  analyzed,  it  is  necessary  first  to  destroy  the  organic  material. 
Several  methods  have  been  proposed: 

(1)  Method  of  Fresenius  and  Babo.— The  tissue  is  cut  into  small  pieces  and 
about  an  equal  weight  of  pure  hydrochloric  acid  added  to  this,  enough  water 
should  be  added  to  form  a  thin  paste  and  dilute  the  hydrochloric  acid  five  or 
six  times.  The  mass  is  heated  on  the  water  bath  and  crystals  of  potassium 
chlorate  added  in  small  amounts  at  a  time  with  stirring  until  a  clear  yellow 
liquid  is  obtained  containing  a  very  small  amount  of  solid  particles.  Th^ 
heating  is  continued  until  there  is  no  odor  of  chlorine,  but  concentration  should 
be  avoided  by  the  addition  of  water.  The  sclu.ion  should  be  cooled  and  filtered, 


§69,  7.  ARSENIC.  69 

the  arsenic  now  being1  present  in  the  filtrate  as  arsenic  acid.  This  solution 
should  be  treated  with  sodium  bisulphite  or  sulphur  dioxide  to  reduce  the 
arsenic  acid  to  arsenous  acid  and  then  the  arsenic  may  be  precipitated  with 
hydrogen  sulphide.  It  is  advisable  to  pass  the  hydrogen  sulphide  through  the 
warm  liquid  for  twenty-four  hours  to  insure  complete  precipitation.  A  yel- 
lowish precipitate  of  organic  matter  will  usually  be  obtained  even  if  arsenic 
be  absent.  The  precipitate  should  be  filtered,  washed,  and  then  dissolved  in 
dilute  ammonium  hydroxide,  which  separates  it  from  other  sulphides  of  the 
silver,  tin  and  copper  groups,  that  may  be  present.  A  portion  at  least  of  the 
precipitated  organic  matter  will  dissolve  in  the  ammonium  hydroxide.  The 
filtrate  should  be  acidulated  with  hydrochloric  acid,  filtered  and  washed. 
Dissolve  the  precipitate  in  concentrated  nitric  acid  and  evaporate  to  dryness. 
Kedissolve  in  a  small  amount  of  water,  add  a  drop  of  nitric  acid,  filter  and  test 
the  filtrate  by  Marsh's  test  or  any  of  the  other  tests  for  arsenic. 

(2)  Hydrochloric  acid  diluted  alone  may  be  used  for  the  disintegration  of 
the  soft  animal  tissues.     The  solution  will  usually  be  dark  colored  and  viscous 
and  not  at  all  suited  for  further  treatment  with  hydrogen  sulphide;  but  may 
be  at  once  subjected  to  the  Reinseh  test  (6'd). 

(3)  Method  of  Danger  and  Flandin. — The  tissue  may  be  destroyed  by  heat- 
ing in  a  porcelain  dish  with  about  one-fourth  its  weight  of  concentrated  sul- 
phuric   acid.     When    the    mass  becomes    dry    and    carbonaceous    it    is    cooled, 
treated  with  concentrated  nitric  acid  and  evaporated  to  dryness.     Moisten  with 
water,  add  nitric  acid,  and  again  evaporate  to  dryness;  and  repeat  until  the 
mass  is  colorle'ss.     Dissolve  in  a  small  amount  of  water  and  test  for  arsenic  by 
the  usual  tests.     This  method  is  objectionable  if  chlorides  are  present  as  the 
volatile  arsenous  chloride  will  be  formed. 

(4)  Method  by  distillation  with  hydrochloric  acid.     The  finely  divided  tissue 
is  treated,  in  a  retort,  with  its  own  weight  of  concentrated  hydrochloric  acid 
and  distilled  on  the  sand  bath.     Salt  and  sulphuric  acid  may  be  used  instead  of 
hydrochloric  acid.     A  receiver  containing  a  small  amount  of  water  is  connected 
to  the  retort  and  the  mass  distilled  nearly  to  dryness.     If  preferred,  gaseous 
hydrochloric  acid  IT  ay  be  conducted  into  the  retort  during  the  process  of  dis- 
tillation, in  which  case  all  the  arsenic  (even  from  arsenous  sulphide  (5c))  will 
be  carried  over  in  the  first  100  cc.  of  the  distillate.     The  receiver  contains  the 
arsenic,  a  great  excess  of  hj'drochloric  acid  and  a  small  amount  of  organic 
matter.     To  a  portion  of  this  solution  the  Reinsch  test  may  be  applied  at  once 
and  other  portions  may  be  diluted  and  tested  with  hydrogen  sulphide  or  the 
solution  may  at  once  be  tested  in  the  Marsh  apparatus. 

For  more  detailed  instructions  concerning  the  detection  and  estimation  of 
arsenic  in  organic  matter,  special  works  on  Toxicology  and  Legal  Medicine 
must  be  consulted.  The  following  are  valuable  works  on  this  subject:  Medical 
Jurisprudence. —  Forensic  Medicine  and  Toxicology,  Witthaus  and  Becker,  Vol. 
iv,  1911;  Laboratory  Manual  for  the  Detection  of  Poisons  and  Powerful  Drugs, 
Dr.  Wilhelm  Autenreith  (translated  by  Wm.  H.  Warren),  1915;  Allen's  Commer- 
cial Organic  Analysis,  Vol.  vi  (4th  edition,  1912);  The  Qualitative  Analysis  of 
Medicinal  Preparations,  H.  C  Fuller,  1912;  Elementary  Chemical  Micros- 
copy, Emile  M.  Chamot,  1915;  Biochemisches  Handlexikon,  V.  Band,  Dr. 
Emil  Abderhalden,  1911;  Mikrochemische  Analyse,  1  and  11  Teil,  P.  D.  C.  Kley, 
1915;  Medical  Jurisprudence,  Taylor;  Ermittelung  von  Giften,  Dragendorff. 

7.  Ignition. — Metallic  arsenic  is  obtained  by  igniting  any  compound 
containing  arsenic  with  potassium  carbonate  and  charcoal,*  or  with  potas- 
sium cyanide : 

2As20,  +  6KCN  =  As,  +  6KCNO 

2As2S3  +  6KCN  =  As4  +  6KCNS 

2As2S8  +  6Na2C03  +  6KCN  =  As4  +  6Na3S  +  6KCNO  +  6CO,  . 

4H,AsO<  +  50  =  As4  +  5C02  +  6H2O 

*  A  very  suitable  carbon  for  the  reduction  of  arsenic  is  obtained  by  igniting  an  alkali  tartrate 
In  absence  of  air  to  complete  carbonization. 


70  A&8ENM.  §69, 8. 

If  this  ignition  be  performed  in  a  small  reduction-tube  *  (a  hard  glass  tube 
about  7  mm.  in  diameter,,  drawn  out  and  sealed  at  one  end),  the  reduced 
arsenic  sublimes  and  condenses  as  a  mirror  in  the  cool  part  of  the  tube. 
The  test  may  be  performed  in  the  presence  of  mercury  compounds,  but 
more  conveniently  after  their  removal;  in  presence  of  organic  material,  it 
is  altogether  unreliable.  If  much  free  sulphur  be  present  the  arsenic 
should  be  removed  by  oxidation  to  arsenic  acid  by  nitric  acid  or  hydro- 
chloric acid  and  potassium  chlorate,  then  precipitation  after  addition  of 
ammonium  hydroxide  by  magnesium  mixture  and  thoroughly  drying  before 
mixing  with  the  cyanide  or  other  reducing  agent. 

8.  Detection. — Arsenic  is  precipitated,  from  the  solution  acidulated  with 
hydrochloric  acid,  in  the  second  group  by  hydrosulphuric  acid  as  the 
sulphide  (6e).     B}r  its  solution  in  (yellow)  ammonium  sulphide  it  is  sepa- 
rated from  Hg,  Pb ,  Bi,  Cu ,  and  Cd  .     By  reduction  to  arsine  in  the 
Marsh  apparatus  it  is  separated  with  antimony  from  the  remaining  second 
group  metals.     The  decomposition  of  the  arsine  and  stibine  with  silver 
nitrate  precipitates  the  antimony,  thus  effecting  a  separation  from  the 
arsenic,  which  passes  into  solution  as  arsenous  acid.     The  excess  of  AgN03 
is  removed  by  HC1  or  CaCl2  and  the  presence  of  arsenic  confirmed  by  its 
precipitation  with  H2S .     For  other  methods  of  detection  consult  the  text 
(6,  6'  and  7).    For  distinction  between  Asv  and  As'"  see  (6  and  §88,  4). 

9.  Estimation. — (1).  As  lead  arsenate,  Pb,(As04)2 .   To  a  weighed  portion 
of  the  solution  containing  arsenic  acid,  a  weighed  amount  of  PbO  is  added. 
After  evaporation  and  ignition  at  a  dull  red  heat  the  residue  is  weighed 
as     Pb3(As04)2  from  which  the  weight  of  the  added  PbO  is    subtracted. 
The  difference  shows  the  amount  of  arsenic  present  reckoned  as  As205 . 
(2).  It  is  precipitated  by  MgS04  in  presence  of  NH4OH  and  NH4C1 ,  and 
after   drying   at   103°,   weighed   as   MgNH4As04.H20  ;   antimony   is   not 
precipitated*if  a  tartrate  be  present  (Lesser,  Z.,  1888,  27,  218).     (-5).  The 
MgNH4As04  is  converted  by  ignition  into  Mg2As207 ,  and  weighed.     (4). 
The  solution  of  arsenous  acid  containing  HC1  is  precipitated  by  H2S  . 

*  As  much  of  the  reduction-glass  tubing  contains  arsenic  (?)  Fresenius  (Z.,  2O,  531  and  22,  397) 
recommends  the  following  modification  of  the  above  method :  A  piece  of  reduction  tubing  about 
16  mm.  diameter  and  15  cm.  long  is  drawn  out  to  a  narrow  tube  at  one  end.  The  other  end  of  the 
tube  is  connected  with  a  suitable  apparatus  for  generating  and  drying  carbon  dioxide.  The 
sample  to  be  tested  is  thoroughly  dried  and  mixed  with  the  dry  cyanide  (or  charcoal)  and  car- 
bonate, placed  in  a  small  porcelain  combustion  boat  and  put  in  the  middle  of  the  reduction 
tube.  The  air  is  then  driven  from  the  tube  by  the  dry  carbon  dioxide  and  the  whole  heated 
gently  until  all  moisture  is  expelled.  The  tube  is  then  heated  to  redness  near  the  point  of  con- 
striction and  when  this  is  done  the  boat  is  heated,  gentljr  at  first  to  avoid  spattering  of  the  fus- 
ing mass,  then  to  a  full  redness  till  all  the  arsenic  has  been  driven  out.  During  the  whole  of  the 
experiment  a  gentle  stream  of  carbon  dioxide  is  passed  through  the  tube.  The  arsenic  collects 
as  a  mirror  in  the  narrow  part  of  the  tube  just  beyond  the  heated  portion.  The  small  end  of  the 
tube  may  now  be  sealed,  the  mirror  collected  by  a  gentle  flame,  driven  to  any  desired  portion  of 
the  tube  and  tested  with  the  usual  tests  (6'  c5).  Compounds  of  antimony  when  treated  in  tin? 
way  do  not  give  a  mirror.  As  small  an  amount  as  0.00001  gram  of  As3O3  will  give  a  distinct  mir- 
ror by  this  method. 


§69,  10.  ARSENIC.  71 

The  precipitate  is  separated  from  free  sulphur  by  solution  in  NHjOH  and 
ivprecipitated  with  HC1 .  It  is  then  dried  at  100°  and  weighed  as  As2S3  . 
(-5).  By  precipitation  as  in  (4)  and  removal  of  sulphur  by  washing  with 
CS2 .  Dry  at  100°  and  weigh  as  As2S3 .  (6).  Uranyl  acetate,  in  presence  of 
ammonium  salts,  precipitates  N"H4U02As04  ;  by  ignition  this  is  converted 
into  uranyl  pyroarsenate  (U02)2As.,07 ,  and  weighed  as  such.  (7).  Small 
amounts  may  be  converted  into  the  metallic  arsenic  mirror  by  the  Marsh 
apparatus  and  weighed  or  compared  with  standard  mirrors  (Gooch  and 
Moseley,  C.  N.,  1894,  70,  207).  (8).  As'"  is  converted  into  Asv  by  a 
graduated  solution  of  iodine  in  presence  of  NaHC03  .  The  end  of  the 
reaction  is  shown  by  the  blue  color  imparted  to  starch.  (9).  As"'  is  Oxi- 
dized to  Asv  by  a  graduated  solution  of  K2Cr207 ,  and  the  excess  of 
K,Cr,07  determined  by  a  graduated  solution  of  FeS04 .  (10).  As"'  is  con- 
verted to  Asv  by  a  weighed  quantity  of  K2Cr207  with  HC1 ,  and  the  excess 
of  chlorine  is  determined  by  KI  and  Na2S203  .  (11).  As'"  is  oxidized  to 
Asv  by  a  graduated  solution  of  KMn04  .  The  end  of  the  reaction  is  indi- 
cated by  the  color  of  the  KMn04  .  (12).  Asv  is  reduced  to  As"'  by  a  grad- 
uated solution  of  HI .  The  action  takes  place  in  acid  solutions.  (13).  In 
neutral  solution,  as  arsenate,  add  an  excess  of  standard  AgN03 ,  and  in  an 
aliquot  part  estimate  the  excess  of  AgN03  with  standard  NaCl .  (14).  Dis- 
tillation as  AsCl3  (Piloty  and  Stock,  B.,  1897,  30,  1649;  see  also  6'e  4). 
(15).  The  arsenic  compound  is  converted  into  AsH,  and  this  passed  into  a, 
solution  of  standard  silver  nitrate,  the  excess  of  which  is  estimated  with 
standard  NaCl  or  the  excess  of  AgN03  is  removed  and  the  arsenous  acid 
titrated  as  in  methods  (9)  or  (11).  (16).  Small  amounts  are  determined 
by  conversion  to  AsH3  and  the  stain  produced  on  mercury  bromide  paper 
compared  with  the  stain  produced  by  known  amounts  of  arsenic.  (Seeker 
and  Smith,  U.  8.  Bull.  No.  147,  p.  212-214).  Many  other  methods  have 
been  recommended. 

10.  Oxidation.— As~"'H3  is  oxidized  to  As'"  by  AgN03 ,  H2SO, ,  H2S04 , 
and  HIOS  ;  and  to  Asv  by  KMn04  (Tivoli,  Qazzetta,  1889,  19,  630),  HN02 , 
HN03 ,  Cl  and  Br  (Parsons,  C.  N.,  1877,  35,  235).  As0  is  oxidized  to  As"' 
by  H262  (Clark,  J.  C.,  1893,  63,  886),  HN03 ,  H2S04  hot,  Cl ,  HC10  ,  HC103 , 
Br,  HBrO;?,  HI03 ,  Ag^  (Senderens,  C.  r.,  1887,  104,  175),  and  to  Asv  by 
the  same  reagents  in  excess  except  H2S04  and  Ag',  which  oxidize  to  As"' 
only.  As'"  is  also  oxidized  to  Asv  in  presence  of  acid  by  Pb02 ,  CrVI;  by 
compounds  of  Co,  Ni ,  and  Mn ,  with  more  than  two  bonds;  and  in 
alkaline  mixture  by  Pb02 ,  Hg20 ,  HgO ,  CuO  ,  K2Cr04 ,  K3Fe(CN)0 ,  etc. 
(Mayer,  J.  pr.,  1880  (2),  22,  103).  Arsine  is  oxidized  to  metallic  arsenic  by 
HgCl2  (Magencon  and  Bergeret,  J.  (7.,  1874,  27,  1008),  and  by  As'",  the  As'" 
also  becoming  As0  (Tivoli,  C.  C.,  1887,  1097).  Asv  and  As"'  are  reduced  to 
metallic  arsenic  by  fusion  with  CO ,  with  free  carbon,  or  with  carbon  com- 
bined, as  H2C204 ,  KCT,  etc.  (7).  By  SnCl2  (6*7)  and  H3P02  (Gd)  in  strong 
HC1  solution;  also  with  greater  or  less  completeness  b 
such  as  Cu ,  Cd ,  Zn  ,  Ms ,  etc.  Rideal  (C.  N..  IfiftUf. 


72  ANTIMONY.  §70,  1. 

HC1  solution;  also  with  greater  or  less  completeness  by  some  free  metals, 
such  as  Cu  ,  Cd  ,  Zn  ,  Mg ,  etc.  Rideal,  (C.  N.,  1885,  51,  292)  recommends 
the  use  of  the  copper-iron  wire  couple  for  the  detection  of  small  quantities 
of  arsenic  by  reduction  to  the  elemental  state.  0.00000?5  grams  may  be 
detected.  In  solution  Asv  is  reduced  to  As'"  by  H3PO, ,  H0S ,  H0S03 , 
Na2S203  (6e),  HC1 ,  HBr ,  HI  (G/),  HCNS ,  etc.  Asv  and  As"'  are  reduced 
to  As~'"H3  by  nascent  hydrogen  generated  by  the  action  of  Zn  and  dilute 
H2S04 ,  or,  in  general,  .by  any  metal  and  acid  which  will  give  a  ready 
generation  of  hydrogen,  as  Zn ,  Sn ,  Fe ,  Mg,  etc.,  and  H2S04  and  HC1 
(Draper,  Dingl,  1872,  204,  320).  As'"  is  reduced  to  As-'"H3  by  nascent 
hydrogen  generated  in  alkaline  solution  as,  Al  and  KOH  ,  Zn  and  KOH , 
sodium  amalgam,  etc.  (separation  from  antimony)  (Davy,  Ph.  C.,  1876, 
17,  275;  Johnson,  C.  N.,  1878,  38,  301). 

§70.  Antimony  (Stibium)  Sb  =  120.2.    Valence  three  and  five  (§11). 

1.  Properties. — Specific  Gravity,  6.62  (Z.  anorgan.  Chem.,  1902,  177).    Melting 
point,  630°  (Cir.  B.  S.,  35,  1915).     Boiling  point,  between  1500°  and  1700°  (B., 
1889,  725).     Its  molecular  weight  is  unknown,  as  its  vapor  density  has  not  been 
taken.     Antimony  is  a  lustrous,   silver  white,   brittle   and   readily  pulverizable 
metal.     It  is  but  little  tarnished  in  dry  air  and  oxidizes  slowly  in  moist  air,  forming 
a  blackish  gray  mixture  of  antimony  and  antimonous  oxide.     At  a  red  heat  it 
burns  in  the  air  or  in  oxygen  with  incandescence,  forming  white  inodorous  (dis- 
tinction from  arsenic)  vapors  of  antimonous  oxide. 

2.  Occurrence. — Native   in    considerable    quantities    in    northern    Queensland, 
Australia  (Mac  Ivor,  C.  N.,  1888,  57,  64);  as  stibnite,  Sb2S3;  as  valentinite,  Sb2O3; 
in  very  many  minerals  usually  combined  with  other  metals  as  a  double  sulphide 
(Campbell,  Pliil.  Mag.,  1860,  (4),  20,  304;  21,  318). 

3.  Preparation. —  (a)  The  sulphide  is  converted  into  the  oxide  by  roasting  in 
the  air,  and  then  reduced  by  fusion  with  coal  or  charcoal.     (6)  The  sulphide  is 
fused  with  charcoal  and  sodium  carbonate:  2Sb2S3  +  6Na2CO3  +30  =  4Sb  + 
6Na2S  +  9CO,  .     (c)  It  is  reduced  by  metallic  iron:  Sb2S3  +  3Fe  =  2Sb  +  3FeS  . 
(d)    To  separate  it  from  other  metals  with  which  it  is   frequently  combined 
requires  a  special  process  according  to  the  nature  of  the  ore  (Dexter,  J.  pr., 
1839,  18,  449;  Pfeifer,  A.,  1881,  209,  161). 

4.  Oxides.— Antimony  forms   three   oxides,    Sb2O3  ,    Sb204  ,    and   Sb,O5  .     (a) 
Antimonous  oxide,  Sb2Os  ,  is  formed  (1)  by  the  action  of  dilute  nitric  acid  upon 
Sb°;  (2)  by  precipitating  SbCls  with  Na2C03  or  NH4OH;  (3)  by  dissolving  Sb° 
in  concentrated  H2S04  and  precipitating  with  Na2C03;  (4)  by  burning  antimony 
at  a  red  heat  in  air  or  oxygen;  (5)  by  heating  Sb2O4  or  Sb205  to  800°  (Baubigny, 
C.  r.,  1897,  124,  499,  and  560).     It  is  a  white  powder,  turning  yellow  upon  heat- 
ing and  white  again  upon  cooling;  melts  at  a  full  red  heat,  becoming  crystalline 
upon  cooling;  slightly  soluble  in  water,  fairly  soluble  in  glycerine  (5&).     Anti- 
monous oxide  sometimes  acts  as  an  acid,  Sb2O3  +  2Na.OH  =  2NaSb02  +  H,O; 
but  more  commonly  as  a  base.     Ortho  and  pyro  antimonous  acids  are  known 
in  the  free  state.     The  meta  compound  exists  only  in  its  salts   (D.,  2,  1,  198). 
(6)   Diantimony  tetroxide,   Sb2O4  ,  is  formed  by  heating  Sb°  ,  Sb2S3  ,   Sb,0,  , 
or  Sb205  in  the  air  at  a  dull  red  heat  for  a  long  time.     The  antimony  in  this 
compound  is  probably  not  a  tetrad,   but  a   chemical  union   of  the   triad   and 
pentad:  2Sb3O4  =  2Sb"'SbvO4  =  Sb2O3.Sb2O5  .     It  is  found  native  as  antimony 
ochre,     (c)    Antimonic   oxide,    Sb,Os  ,    is    formed    by    treating    Sb°  ,    Sb,O3    or 
Sb204   with  concentrated   nitric  acid.     When  heated   to   300°   it  loses  oxygen, 
forming  Sb204  (Geuther,  J.  pr.,  1871,  (2),  4,  438).     It  is  a  citron-yellow  powder, 
insoluble  in  water  but  reddening   moist  blue   litmus   paper.     Antimonic   acid 
exists  in  the  three  *  forms,   analogous  to  the   arsenic   and  phosphoric   acids, 

*Beilstein  and  Blaese  (C.  C.,  1889,  803'  have  prepared  a  number  of  antimonates  and  conclude 
that  the  acid  is  always  the  meta,  H  SbO, . 


£70,  oft.  ANTIMONT.  73 

j.  c.,  ortho,  meta  and  pyro  (Geuther,  I.  c.,  and  Conrad,  C.  N.,  1879,  40,  198).  The 
ortho  acid,  H3SbO4  is  formed  by  the  decomposition  of  the  pentachloride  with 
water  and  washing-  until  the  chloride  is  all  removed  (Conrad,  I.  c.,  and  Dau- 
brawa,  A.,  1877,  186,  110).  The  most  of  the  antimonates  formed  in  the  wet  way 
by  precipitation  from  the  acid  solution  of  antimonic  chloride  are  the  ortho 
antimonates.  By  heating-  the  ortho  acid  to  200°  the  meta  acid,  HSb03  ,  is 
formed.  Strong  ignition  of  Sb,O3  with  potassium  nitrate  and  extraction  with 
water  gives  the  potassium  metantimonate,  KSbOs  ,  and  by  adding  nitric  acid 
to  a  solution  of  this  salt  the  free  acid  is  formed.  The  ortho  acid  dried  at  100° 
gives  the  pyro  acid:  2H3SbO4  =  H4Sb2O7  -f  H,O  (Conrad,  L  c.),  which  upon 
further  heating  to  200°  gives  the  meta  acid.  The  pyroantimonic  acid  forms 
two  series  of  salts,  M4Sb207  and  M2H2Sb207  .  The  sodium  salt  Na2H2Sb2OT 
is  insoluble  in  water  and  is  formed  in  the  quantitative  estimation  of  antimony 
(9),  and  also  in  a  method  for  the  detection  of  sodium  (§206,  6*7).  For  the  latter 
the  soluble  potassium  salt  K2:Er,Sb207  is  used  as  the  reagent.  It  is  prepared 
by  fusing  antimonic  acid  with  a  large  excess  of  potassium  hydroxide;  then 
dissolving,  filtering,  evaporating  and  digesting  hot,  in  syrupy  solution,  with  a 
large  excess  of  potassium  hydroxide,  best  in  a  silver  dish,  decanting  the 
alkaline  liquor,  and  stirring  the  residue  to  granulate,  dry.  This  reagent  must 
be  kept  dry,  and  dissolved  when  required  for  use;  inasmuch  as,  in  solution,  it 
changes  to  the  tetrapotassium  pyroantimonate,  K4Sb2O7  ,  which  does  not 
precipitate  sodium.  The  reagent  is,  of  course,  not  applicable  in  acid  solutions. 
The  reaction  is  as  follows:  K2H2Sb2O7  +  2NaCl  =  Na2H2Sb207  +  2KC1  (§11). 

The  ortho  acid,  HsSb04  ,  is  sparingly  soluble  in  water,  easily  soluble  in  KOH, 
but  insoluble  in  NaOH.  The  meta  acid,  HSbO8  ,  is  sparingly  soluble  in  water, 
easily  soluble  in  both  the  fixed  alkalis;  the  pyro  acid,  H4Sb207  ,  is  sparingly 
(more  easily  than  the  meta)  soluble  in  water;  the  normal  fixed  alkali  salts, 
R4SbnO7  ,  are  soluble  in  water,  also  the  acid  potassium  salt,  K2H2Sb207  ,  but 
not  the  corresponding  sodium  salt,  Na2H._,Sb2O7  . 

5.  Solubilities.— a. — Metal. — Antimony  is  attacked  but  not  dissolved  by  nitric 
acid,  forming  Sb2O3  (a)  or  Sb2O5  (&),  depending  upon  the  amount  and  degree 
of  concentration  of  the  acid;  it  is  slowly  dissolved  by  hot  concentrated  sulphuric 
acid,  evolving  S02  and  forming  Sb2(S04)3  (c) ;  it  is  insoluble  in  HC1  out  of  con- 
tact with  the  air,  but  the  presence  of  moist  air  causes  the  oxidation  of  a  small 
amount  of  the  metal  to  Sb,O3  ,  which  is  dissolved  in  the  acid  without  evolution 
of  hydrogen  (Ditte  and  Metzner,  A.  Ch.,  1896,  (6),  29,  389). 

The  best  solvent  for  antimony  is  nitric  acid,  followed  by  hydrochloric  acid  or 
nitrohydrochloric  acid  containing  only  a  small  amount  of  nitric  acid.  Anti- 
monous  chloride,  SbCl3  ,  is  at  first  formed  (d),  but  if  sufficient  nitric  acid  be 
present  this  is  rapidly  changed  to  antimonic  chloride,  SbCl5  (e).  If,  however, 
too  much  nitric  acid  be  present,  the  corresponding  oxides  (not  readily  soluble 
in  nitric  acid)  are  precipitated  (6c).  The  halogens  readily  attack  the  metal 
forming  at  first  the  corresponding  trihalogen  compounds  (d).  Chlorine  and 
bromine  (gas)  unite  with  the  production  of  light,  and  if  the  halogen  be  in 
excess,  the  pentad  chloride  (e)  or  bromide  is  formed  (Berthelot  and  Petit,  A.  Ch., 
1891,  (6),  18,  65).  The  pentiodide,  SbI5  ,  does  not  appear  to  exist  (Mac  Ivor, 
J.  C.,  1876,  29,  328). 

(a)  2Sb  +  2HN03  =  Sb203  +  2ND  +  H2O 

(6)  6Sb  +  10HNO3  =  3Sb2O5  -f  10NO  +  5HZ0 

(c)  2Sb  +  f,H2SO4  =  Sb2(SO4)3  -f  3S02  +  6H2O 

(d)  2Sb  +  3CL  =  2SbCls 

(e)  SbCl3  +  C12  =  SbCl5 

ft. — Oaridett. — Antimonous  oxide,  Sb2O3  ,  is  soluble  in  55,000  parts  of  water  at 
15°  and  in  10,000  parts  at  100°  (Schulze,  J.  Pr.,  1883,  (2),  27,  320);  insoluble  in 
alcohol;  soluble  in  hydrochloric  (a),  sulphuric  and  tartaric  (&)  acids  with 
formation  of  the  corresponding  salts.  The  dry  ignited  oxide  is  scarcely  at  all 
soluble  in  nitric  acid;  the  moist,  freshly  precipitated  oxide,  on  the  other  hand, 
dissolves  readily  in  the  dilute  or  concentrated  acid,  be  it  hot  or  cold.  Under 
certain  conditions  of  concentration  a  portion  of  the  antimony  precipitates  out 
upon  standing  as  a  white  crystalline  precipitate.  It  is  soluble  in  the  fixed 


74  A\TLMOXY.  §70,  5c. 

alkali  hydroxides  with  formation  of  metantimonites  (c)  (Terreil,  A.  Cli.,  1866, 
(4),  7,  350).  Fixed  alkali  carbonates  dissolve  a  small  amount  of  the  oxide  with 
the  probable  formation  of  some  antimonite  (rf)  (Schneider,  Poyy.,  1859,  108,  407). 
It  is  fairly  soluble  in  glycerine  (Kohler.  DingL,  18S5.  258.  520). 

(a)  Sb2O3  +  6HC1  =  2SbCl3  +  3H2O 

(6)  Sb203  +  H2C4H40G  =  (SbO)2C4H400  .+  H20 

(c)  Sb203  +  2KOH  =  2KSb02  +  H20 

(d)  Sb203  +  NasCOs  =  2NaSb02  +  CO2 


Antimony  tetroxide,  Sb204  ,  is  insoluble  in  water,  slowly  dissolved  by  hot 
concentrated  hydrochloric  acid.  Antimonic  oxide,  Sb2O,  ,  is  insoluble  in  water; 
soluble  in  hydrochloric  and  tartaric  acids  without  reduction;  hydriodic  acid 
dissolves  it  as  antimonous  iodide  with  liberation  of  iodine  (Of)  ;  slowly  soluble 
in  concentrated  fixed  alkalis;  soluble  in  alkaline  solution  of  glycerine  (Kohler, 
J.  C.,  1886,  50,  428).  The  hydrated  oxides  of  antimony  (acids)  "have  essentiallv 
the  same  solubilities  as  the  oxides  (4). 

c.  —  Salts.  —  Antimonous  chloride,  SbCL,  ,  is  very  (lelniucwent.  decomposed  by 
pure  water,  forming  a  basic  salt;  soluble  in  water  strongly  acidulated  with  an 
inorganic  acid,  or  tartaric,  citric,  or  oxalic  acids  (6&),  but  not  when  acidulated 
with  acetic  acid;  it  is  also  soluble  in  concentrated  solutions  of  the  chlorides  of 
the  alkalis  and  of  the  alkaline  earths  (Atkinson,  (''.  .V.,  1883,  47.  175).  The 
bromide  and  iodide  are  dcl'uiucsccnt  and  require  moderately  concentrated  acid  lo 
keep  them  in  solution.  The  sulphate,  Sb,  (S04)3  ,  dissolves  in  moderately  con- 
centrated sulphuric  acid.  Antimonous  tartrate  and  the  potassium  antimonous 
tartrate  (tartar-emetic)  are  soluble  in  water  without  acidulation;  the  latter  is 
soluble  in  glycerine  and  insoluble  in  alcohol.  The  trichloride,  bromide  and 
iodide  are  soluble  in  hot  CS2;  the  chloride  and  bromide  are  soluble  in  alcohol 
without  decomposition,  but  the  iodide  is  partially  decomposed  by  alcohol  or 
ether  (Mac  Ivor,  J.  (7.,  1876,  99,  328). 

The  pentachloride,  SbCl5  ,  is  a  liquid,  very  readily  combining  with  a  small 
amount  of  w^ater  to  form  crystals  containing  one  or  four  molecules  of  water. 
The  addition  of  more  water  decomposes  the  salt  forming  the  basic  salt;  if, 
however,  a  few  drops  of  HC1  have  been  added  first,  any  desired  amount  of 
water  (if  added  at  one  time)  may  be  added  without  causing  a  precipitation  of 
the  basic  salt.  If  after  acidulation  water  be  added  slowly,  the  basic  salt  will 
soon  be  precipitated. 

Antimonous  sulphide,  Sb2S3  ,  is  readily  soluble  in  K2S  ,  and  on  evapora- 
tion large  yellow  transparent  crystals  of  K4Sb2S.  are^obtained  (a)  (Ditte, 
C.  r.,  1886,  102,  1G8  and  212).  It  is  soluble  in  moderately  concentrated 
HC1  with  evolution  of  ELS  (5);  slowly  decomposed  by  boiling  with  water 
into  Sb203  and  H2S  (c):  and  on  boiling  with  NH4C1  into  SbCl,  and  (NH4),S 
(de  Clermont,  C.  r.,  1879,  88,  972).  Dilute  H2S04  is  almost  without  action, 
dilute  HNO.,  gives  Sb203  (d).  Sparingly  soluble  in  hot  NH4OH  solution, 
soluble  in  the  fixed  alkalis  (on  fusion  or  boiling)  (e);  insoluble  in  (NH4)2C0$ 
(distinction  from  arsenic);  insoluble  in  the  fixed  alkali  carbonates  in  the 
cold  but  on  warming  they  effect  complete  solution  (f)  (distinction  from 
tin);  very  sparingly  soluble  in  normal  ammonium  sulphide;  readily  soluble 
in  yellow  ammonium  sulphide  with  oxidation  (g)  (6e).  The  pentasulphide, 
Sb^S.  ,  is  insoluble  in  water;  soluble  in  the  alkali  sulphides  (Ji),  and  in  the 
fixed  alkali  carbonates  and  hydroxides;  insoluble  in  ammonium  carbonate 
and  sparingly  soluble  in  ammonium  hydroxide,  more  readily  when  warmed 
(D.,  2,  1,  217).  On  boiling  with  water  it  slowly  decomposes  into  Sb,0,., 


§?0,  5d.  ANTIMONV.  75 

H2S  and  S  (Mitscherlich,  J.  pr.,  1840,  19,  455).     Hydrochloric  acid  on 
warming  dissolves  it  a?  SbCl,  (/): 

(a)     Sb2S3  +  2K2S  =  K,Sb,S5 

(1))     Sb2S3  +  C.EC1  =  2SbCL  +  3H2S 

(c)  Sb2S3  +  3H,0  =  Sb203  +  3H2S 

(d)  2Sb2S3  +  4HNO3  =  2Sb,03  +  3S2  +  4NO  +  2H20 

(e)  2Sb,S3  +  4KOH  =  3KSbS2  +  KSbO2  +  2H20 

(/•)  2Sb,S3  -f  2Nsi,CO,  =  :;NaSbS2  +  NaSbO,  +  2CO2 

(y)  2Sb,S3  +  6(NH4)2S2  =  4(NH4)3SbS4  +  S2 

(70  SbaS5  +  3(NH4)2S  =  2(NH4)3SbS4 

(i)  Sb2S5  +  6HC1  =  2SbCl3  +  3H2S  +  S2 

d. — Water.* — With  the  exception  of  the  compounds  of  antimony  with 
some  organic  acids,  as  tartaric  and  citric,  all  salts  of  antimony  are  decom- 
posed by  pure  WATER.  For  this  reason  it  will  be  seen  that  water  is  a  very 
important  reagent  in  the  analysis  of  antimony  salts.  The  salts  with 
inorganic  acids  all  require  the  presence  of  some  free  acid  (not  acetic)  to 
keep  them  in  solution.  If  the  acid  be  tartaric  the  further  addition  of 
water  causes  no  precipitation  of  the  antimony  salt.  Water  decomposes 
the  inorganic  acid  solutions  precipitating  the  basic  salt,  setting  more  acid 
free  which  dissolves  a  portion  of  the  basic  salt.  The  addition  of  more 
water  causes  a  further  precipitation  and  at  the  same  time  dilutes  the  acid 
so  that  upon  the  addition  of  a  sufficient  amount  of  water  a  nearly  com- 
plete precipitation  may  be  obtained.  If  the  precipitate  of  the  basic  salt  be 
washed  with  water  the  acid  is  gradually  displaced,  leaving  finally  the  anti- 
mony as  oxide. 

With  solutions  of  antimonous  chloride  the  basic  salt  precipitated  is 
white  antimonous  oxychloride,  Sb4Cl205 ,  "  Powder  of  Algaroth,"  soluble 
in  tartaric  acid  (distinction  from  bismuth,  §76,  5d)  (Mac  Ivor,  C.  N.,  1875, 
32,  229),  4SbCl3  +  5H,0  =  Sb4Cl,05  +  10HC1 .  The  basic  salt  repeatedly 
washed  with  water  is  slowly  (rapidly  if  alkali  carbonate  be  used)  changed 
to  the  oxide,  Sb203  (Malaguti,  J.  pr.,  1835,  6,  253),  Sb4Cl205  +  H20  = 
2Sb,03  +  2HC1.  With  antimonic  chloride,  SbCl5  ,  the  basic  salt  is 
SbOCl, ;  SbCl,  +  H20  =  SbOCl3  +  2HC1  (Williams,  (7.  N.,  1871,  24,  224). 

Solutions  of  the  tartrates  of  antimony  and  of  antimony  and  potassium 
are  not  precipitated  on  the  addition  of  water;  and  antimonous  chloride 

*The  acidity  of  water  solutions  of  certain  salts  having-  a  weak  base  and  the  alkalinity  of 
others  containing  a  weak  acid  is  due  to  a  partial  decomposition  (hydrolysis)  of  the  salt  by  the 
ions  of  the  water,  H*  and  OH',  forming-  again  the  original  acid  and  base.  ]Va2CO3,  for  instance, 
is  split  up  into  the  weak  non-dissociated  H2CO3  and  the  strongly-dissociated  NaOH,  whose 
OH  ions  give  the  "alkaline  reaction."  FeCl3  in  water  forms  soluble  colloidal  Fe(OH)a,  which 
may  be  separated  by  dialysis  from  the  free  HC1  resulting  or  precipitated  by  addition  of  a 
neutral  salt,  as  NaCl,  to  the  dilute  solution;  KCN  gives  alkaline  KOII  and  non-dissociated 
HCW,  readily  detected  by  its  odor.  In  other  cases  precipitation  is  caused,  as  in  the  treatment 
of  bismuth  or  antimony  solutions  with  water  or  on  heating  "VaoZnOa  solution,  hydrolysis  in 
general  being  increased  by  raising  the  temperature.  The  action  of  water  on  soap  belongs  to 
this  class. 


76  ANTIMONY.  §70,  60. 

dissolved  in  excess  of  tartaric  or  citric  acid  solution  is  not  precipitated  on 
addition  of  water. 

G.  Reactions.— a.— The  alkali  hydroxides  and  carbonates  precipitate  from 
acidulated  solutions  of  inorganic  antimonous  salts,  antimonous  oxide  *  Sb  0 
(a)  (Rose,  Pogg.,  1825,  3,  441),  white,  bulky,  readily  becoming-  crystalline  on 
boiling;  sparingly  soluble  in  water  (56),  readily  soluble  in  excess  of  the  fixed 
alkalis,  forming  a  metantimonite  (6)  (Terreil,  A.  Ch.,  1866,  (4),  7,  350);  slowly 
soluble  in  a  strong  excess  of  a  hot  solution  of  the  fixed  alkali  carbonate  (c) 
(distinction  from  tin);  insoluble  in  ammonium  hydroxide  or  ammonium  car- 
bonate. The  freshly  precipitated  oxide  is  readily  soluble  in  acids  (not  in  acetic 
acid).  If  the  alkaline  solution  of  the  antimony  be  carefully  neutralized  with 
an  acid  (not  tartaric  or  citric)  the  oxide  is  precipitated  (d)  and  at  once  dissolved 
by  further  addition  of  acid.  The  presence  of  tartaric  or  citric  acids  prevents 
the  precipitation  of  the  oxide  by  means  of  the  alkalis  or  alkali  carbonates. 

Antimonous  oxide  acts  as  a  feebly  acidic  anyhdride  toward  alkalis,  with 
which  it  combines,  dissolving  in  their  solutions  and  forming  antimonites,  which 
are  found  to  be  monobasic,  so  far  as  capable  of  isolation.  Sodium  ar.timonite, 
NaSbO,  ,  is  the  most  stable  and  the  least  soluble  in  water;  potassium  anti- 
monite,  KSbO2  ,  is  freely  soluble  in  dilute  potassium  hydroxide  solution,  but 
decomposed  by  pure  water.  By  long  standing  (24  hours),  a  portion  of  the 
antimonous  oxide  deposits  from  the  alkaline  solution,  and  the  presence  of  alkali 
hydrogen  carbonates  causes  a  nearly  complete  separation  of  that  oxide  (e). 
(a)  2SbCl3  +  6KOH  =  Sb203  +  6KC1  +  3H20 

2SbCls  +  3Na2C03  =  Sb203  +  6NaCl  +  3CO2 
(6)     Sb20,  +  2KOH  =  2KSb02  +  H20 
or  SbCl3  +  4KOH  =  KSb02  +  3KC1  +  2H2O 

(c)  Sb2O3  +  Na,CO3  =  2NaSb02  +  C02 

(d)  2KSb02  +  2HC1  =  Sb203  +  2KC1  +  H2O 

(e)  2NaSb02  +  2NaHC03  =  Sb208  +  2Na2C03  +  H20 

Antimonic  salts  are  precipitated  under  the  same  conditions  as  the  antimonous 
salts.  The  freshly  formed  precipitate  is  the  orthoantimonic  acid,  HsSb04  = 
SbO(OH)8  =  Sb2O5,3H2O  (a)  (Conrad,  C.  N.,  1879,  40,  198);  insoluble  in  am- 
monium hydroxide  or  carbonate;  soluble,  more  readily  upon  warming,  in 
excess  of  the  fixed  alkali  hydroxides  and  carbonates  as  metantimonate  (6). 

(a)     SbCl5  +  5KOH  =  Sbo'(OH),  +  5KC1  +  H2O 
(6)     SbO(OH)8  +  KOH  =  KSb03  +  2H2O 

6. — The  freshly  precipitated  antimonous  oxide  is  soluble  in  oxalic  acid,  but 
(in  absence  of  tartaric  acid)  the  antimony  soon  slowly  but  completely  separates 
out  as  a  white  crystalline  precipitate;  unless  an  alkali  oxalate  be  present,  when 
the  soluble  double  oxalate  is  formed.  The  precipitate  of  antimony  oxalate 
dissolves  upon  the  further  addition  of  hj^drochloric  acid.  Freshly  precipitated 
antimonic  oxide  dissolves  readily  in  oxalic  acid  and  does  not  separate  out  upon 
standing.  Acetic  acid  precipitates  the  solutions  of  antimony  salts  if  tartaric 
acid  be  absent.  Potassium  cyanide  gives  a  white  precipitate  with  antimonous 
salts  soluble  in  excess  of  the  cyanides. 

With  potassium  ferrocyanide  antimonous  chloride  (not  tartrate)  gives  a 
white  precipitate,  soluble  in  hydrochloric  acid  (distinction  from  tin),  or  fixed 
alkali  hydroxides  (Warren,  C.  N.,  1888,  57,  124).  Potassium  ferricyanide  is 
reduced  to  ferrocyanide  by  antimonous  salts  in  alkaline  solution  (Baumann, 
Z.  angew.,  1892,  117). 

c. — From  the  solutions  of  the  fixed  alkali  antimonites  or  antimonatee  the 
oxides  or  hydrated  oxides  (acids)  are  precipitated  upon  neutralization  with 
nitric  acid  (or  other  inorganic  acids) ;  the  freshly  formed  precipitates  readily 

*  Menschutkin  (page  186)  says  the  precipitate  formed  by  the  action  of  alkalis  upon  antimonous 
salts  is  the  meta  acid,  HSbOa. 


§70,  6«.  ANTIMONY.  77 

dissolving-  in  an  excess  of  the  acid.  Antimonous  nitrate  is  very  unstable  and 
the  antimonic  nitrate  is  not  known  to  exist.  It  is  quite  probable  that  these 
solutions  in  nitric  acid  are  merely  solutions  of  some  of  the  hydrated  oxides 
(acids). 

d.   Compounds  of  antimony  with  the   acids  of   phosphorus  are  not   known, 
(Na2HP04  does  not  precipitate  antimony  salts,  separation  from  tin,  §71,  Qd). 

c.  Hydrogen  sulphide  precipitates,  from  acid  *  solutions  of  antimonous 
salts,  antimonous  sulphide  (a),  Sb2S3 ,  orange-red;  in  neutral  solutions 
(tartrates)  the  precipitation  is  incomplete.  In  strong  fixed  alkali  solu- 
tions (6a)  the  precipitation  is  prevented,  or  rather  the  sulphide  first 
formed  (&)  is  at  once  dissolved  in  the  excess  of  the  fixed  alkali  (c),  sparingly 
in  NH4OH .  The  alkali  sulphides  give  the  same  precipitate  sparingly 
soluble  in  normal  ammonium  sulphide,  readily  soluble  in  the  fixed  alkali 
sulphides  (d)  and  in  yellow  ammonium  sulphide  (e).  Antimonous  sulphide 
is  slowly  decomposed  by  boiling  water  (f) ;  insoluble  in  ammonium  carbon- 
ate (distinction  from  As);  slowly  soluble  in  boiling  solution  of  the  fixed 
alkali  carbonates  (g)  (distinction  from  Sn) ;  soluble  in  hot  moderately  con- 
centrated hydrochloric  acid  (h)  (distinction  from  arsenic).  The  alkaline 
solutions  of  antimonous  sulphide  are  oxidized  upon  standing  by  the  oxygen 
of  the  air  or  rapidly  in  the  presence  of  sulphur  (e) ;  from  the  alkaline  solu- 
tions hydrochloric  acid  precipitates  the  antimony  as  trisulphide,  penta 
sulphide  or  a  mixture  of  these,  depending  upon  the  degree  of  .oxidation  (i). 
(a)  2SbCls  +  3H2S  =  Sb2S3  +  6HC1 
(6)  2KSb02  +  3H2S  =  Sb2S3  +  2KOH  +  2H2O 

(c)  2Sb2S3  +  4KOH  =  3KSbS2  +  KSbO2  +  2H2O 

(d)  Sb2S3  +  K2S  =  2KSbS2 

(e)  2Sb2S3  +  6(NH4)2S2  =  4(NH4)3SbS4  +  Sa 

(f)  Sb2S3  -f  3H20  =  Sb20s  +  3H2S 

(g)  2Sb2S3  +  2K2C03  =  3KSbS2  -f-  KSbO2  +  2C02 
(ft)     Sb2S3  +  6HC1  =  2SbCls  +  3H2S 

(i)      3KSbS2  -f  KSb02  +  4HC1  =  2Sb2S3  +  4KC1  +  2H20 
or  2(NH4)3SbS4  +  6HC1  =  Sb2S6  +  6NH4C1  +  3H2S 

Hydrosulphuric  acid  f  and  alkali  sulphides  precipitate  (under  like  condi- 
tions as  for  antimonous  salts),  from  solutions  of  antimonic  salts,  antimonic 
sulphide,  Sb2S5 ,  orange,  having  the  same  solubilities  as  the  tri-sulphide. 
The  alkaline  solution  of  the  sulphide  consists  chiefly  of  the  ortho-thioanti- 
monate  instead  of^the  meta,  as  in  antiinpnous  compounds.  Sb2S5  -f  3K2S 
=  2K3SbS4  ;  4Sb2S5  +  18KOH  =  5K3SbS4  +  3KSb03  +  9H20  .  When 
dissolved  in  HC1  the  penta-sulphide  is  reduced  to  SbCl3  with  liberation 
of  sulphur,  Sb2S5  +  6HC1  =  2SbCl3  +  3H2S  +  S2 . 

*  According  to  Loviton  (J.  C.,  1888,  54,  993)  the  precipitation  takes  place  in  the  presence  of 
quite  strong  hydrochloric  acid  tone  to  one)  separation  from  tin,  which  is  precipitated  only  when 
three  or  more  parts  of  water  are  present  to  one  of  the  acid.  See  also  Noyes  and  Bray,  J.  Am. 
Soc.t  29,  137  (1917). 

t  In  order  to  precipitate  pure  antimonic  sulphide,  the  solution  of  the  antimonic  salt  must  be 
cold,  and  the  hydrogen  sulphide  added  rapidly.  If  the  solution  be  warmed  or  the  hydrogen  sul- 
phide added  slowly  more  or  less  antimonous  sulphide  is  precipitated  (B6sek,.7.  C.,  1895,  67,515). 


78  ANTIMONY.   *  §70,  6/. 

All  salts  of  antimony  when  warmed  with  sodium,  thiosulphate,  Na2S2O3  , 
are  precipitated  as  the  sulphide  (separation  of  arsenic  and  antimony).  2SbCl3 
+  3Na2S203  +  3H2O  ==  Sb2S3  +  3Na,,SO4  +  GHC1 .  Sulphurous  acid  reduces 
antimonic  salts  to  antimonous  salts  (Knorre,  Z.  angeic.,  1888,  155).  Sulphates  of 
antimony  are  not  prepared  by  precipitation,  but  by  boiling  the  oxides  with 
strong  sulphuric  acid.  They  dissolve  only  in  very  strongly  acidulated  water. 

/. — Antimony  occurs  most  frequently  for  analysis  as  the  chlorides;  it  is 
therefore  important  that  the  student  familiarize  himself  with  the  deport- 
ment of  these  salts  with  the  various  reagents,  used  in  qualitative  analysis. 
The  most  important  of  the  properties  have  been  discussed  under  5a,  &.  c,  d. 
Hydrochloric  acid,  or  any  other  inorganic  acid,  carefully  added  to  a  solu- 
tion of  antimony  salts  in  the  fixed  alkalis  will  precipitate  the  correspond- 
ing oxide  or  hydrated  oxide,  soluble  upon  further  addition  of  the  acid. 
Potassium  iodide  added  to  antimonous  chloride  solution,  not  too  strongly 
acid,  gives  a  yellow  precipitate  of  antimonous  iodide,  soluble  in  hydro- 
chloric acid.  The  precipitation  does  not  take  place  in  the  presence  of 
tartaric  or  oxalic  acids.  Hydriodic  acid  (or  potassium  iodide  in  acidu- 
lated solutions)  added  to  solutions  of  antimonic  salts  causes  a  reduction 
of  the  antimony  to  an  antimonous  salt  with  liberation  of  iodine  (distinc- 
tion from  SnIV:  SbCl5  +  2HI  =  SbCl3  +  2HC1  +  I2 .  The  iodine  may  be 
detected  by  heating  and  obtaining  the  violet  vapors,  or  by  adding  carbon 
disulphide  and  shaking.  It  should  be  remembered  that  the  solution  to 
be  tested  must  be  acid,  for  in  alkaline  solutions  the  reverse  action  takes 
place,  iodine  oxidizing  antimonous  salts  to  antimonic  salts:  SbCl3  -)- 
8KOH  +  I2  —  K3Sb04  +  2KI  +  3KC1  +  4H20  (Weller,  A.,  1882,  213, 
364).  Also  the  absence  of  other  oxidizing  agents  which  liberate  iodine 
from  hydriodic  acid  must  be  assured. 

g. — If  antimony  and  arsenic  compounds  occurring  together  are  strongly 
oxidized  with  nitric  acid  there  is  danger  that  the  insoluble  precipitate  of  anti- 
monic oxide  may  contain  arsenic,  as  antimonic  arsenate,  insoluble  (Menschut- 
kin).  Stannous  chloride  reduces  antimonic  compounds  to  the  antimonous 
condition,  but  in  no  case  causes  a  precipitation  of  the  metal  (distinction  from 
arsenic). 

Ji. — Antimonous  salts  in  acid,  neutral  or  alkaline  solution,  rapidly  reduce 
solutions  of  chromates  to  chromic  compounds.  Acid  solutions  of  antimonous 
salts  reduce  solutions  of  manganates  and  permanganates  to  manganous  salts; 
with  alkaline  solutions  to  manganese  dioxide.  These  reactions  are  capable  of 
quantitative  application  in  absence  of  other  reducing  agents.  The  antimony  is 
oxidized  to  the  antimonic  condition  (9  and  10). 

i. — An  antimonous  compound  when  evaporated  on  a  water  bath  with  an 
ammoniacal  solution  of  silver  nitrate  gives  a  black  precipitate  (Bun sen,  A., 
1855,  106,  1).  A  solution  of  an  antimonous  compound  in  fixed  alkali  when 
treated  with  a  solution  of  silver  nitrate  gives  a  heavy  black  precipitate  of 
metallic  silver,  insoluble  in  ammonium  hydroxide,  and  thus  separated  from  the 
precipitated  silver  oxide.  If  instead  of 'a  water  solution  of  silver  nitrate,  a 
solution  with  great  excess  of  ammonium  hydroxide  (one  to  sixteen)  be  added, 
no  precipitation  occurs  in  the  cold  (distinction  from  Sn") ;  nor  upon  heating 
until  the  excess  of  ammonia  has  been  driven  off.  Antimonates  with  silver 
nitrate  give  a  white  precipitate  of  silver  antimonate,  soluble  in  ammonium 
hydroxide. 


£70,  6/.  ANTIMONY.  79 

j. — Stibine. — By  the  action  of  zinc  and  sulphuric  or  hydrochloric  acid  all 
compounds  of  antimony  are  first  reduced  to  the  metallic  state.  The 
formation  of  stibine  is  a  secondary  reaction  and  requires  the  moderately 
rapid  generation  of  hydrogen  in  acid  solution.  If  a  few  drops  of  a  solu- 
tion of  an  antimony  salt,  acidulated  with  hydrochloric  acid,  be  placed 
upon  a  platinum  foil  and  a  small  piece  of  zinc  be  added,  the  antimony  is 
immediately  deposited  as  a  black  stain  or  coating  adhering  firmly  to  the 
platinum;  2SbCl3  +  3Zn  --  2Sb  -f  3ZnCl, .  In  this  test  tin,  if  present, 
deposits  as  a  loose  spongy  mass,  while  arsenic,  if  present,  does  not  adhere 
so  firmly  to  the  platinum  as  the  antimony.  In  the  presence  of  arsenic 
this  test  should  be  applied  with  caution  under  a  hood  as  a  portion  of  the 
arsenic  is  almost  immediately  evolved  as  arsine  (§69,  6'&). 

If  hydrogen  be  generated  more  abundantly  than  in  the  operation  above 
mentioned,  by  zinc  and  dilute  sulphuric  or  hydrochloric  acid,  the  gaseous 
antimony  hydride,  stibine,  SbH3 ,  is  obtained  for  examination.  For  com- 
parison with  arsine  and  details  of  manipulation  see  "  Marsh's  Test "  under 
arsenic  (§69,  6' a) : 

Sb,O3  +  GZn  +  6H2S04  =  GZnSO,  +  3H20  +  2SbH, 

SbCl3  +  3Zn  +  3HC1  =  SZnCL  +  SbH3 

Stibine  is  a  colorless,  odorless  gas,  not  nearly  so  jpoisonous  as  arsine.  It 
burns  with  a  luminous  and  faintly  bluish-green  flame,  dissipating  vapors 
of  antimonous  oxide  and  of  water  (a);  or  depositing  antimony  on  cold 
porcelain  held  in  the  flame,  as  a  lusterless  brownish-black  spot  (&).  The 
gas  is  also  decomposed  by  passing  through  a  small  glass  tube  heated  to 
low  redness  (c),  forming  a  lustrous  ring  or  mirror  in  the  tube.  The  stibine 
is  decomposed  more  readily  by  heat  than  the  arsine  and  the  mirror  is 
deposited  on  both  sides  of  the  heated  portion  of  the  glass  tube.  The  spots 
and  mirror  of  antimony  are  compared  with  those  of  arsenic  in  §69,  6'c. 
The  antimony  in  stibine  is  deposited  as  the  metal  when  the  gas  is  passed 
into  a  concentrated  solution  of  fixed  alkali  hydroxide  or  when  it  is  passed 
through  a  IT  tube  filled  with  solid  caustic  potash  or  soda-lime  (distinction 
and  separation  from  arsenic). 

(a)     2SbH3  +  302  =  Sb,03  +  3H2O 

(6)      4SbH3  -f  302  =  4Sb  +  6H2O 

(c)     2SbH3  =  2Sb  +  3H2 

When  the  antimony  hydride  (stibine)  is  passed  into  a  solution  of  silver 
nitrate,  the  silver  is  reduced,  leaving  the  antimony  with  the  silver,  as 
(tiiHntnuonx  ur<j<>niide9  SbAg3 ,  a  black  precipitate,  distinction  front  arsenic, 
which  enters  into  solution  (§69,  6'a  and  &);  SbH3  -f  3AgN03  —  SbAg3  + 
3HNO,  .  The  precipitate  should  be  filtered  and  washed  free  from  unde- 
composed  silver  salt  (and  arsenous  acid,  if  that  be  present),  and  dissolved 
\\ith  dilute  hydrochloric  acid  (HC1  does  not  dissolve  uncombined  anti- 


80  ANTiuoyr.  §70, 7. 

mony,  5o) :  SbAg3  +  6HC1  =  SbCL  +  3AgCl  +  3H, .  The  solution  con- 
sists of  antimonous  chloride,,  leaving  silver  chloride  as  a  precipitate. 
However,  in  the  excess  of  hydrochloric  acid  used  a  small  portion  of  the 
silver  chloride  may  be  dissolved  (§59,  oc),  interfering  with  the  final  test 
for  the  antimony.  If  this  be  the  case  the  silver  should  be  removed  by  a 
drop  of  potassium  iodide  (8). 

Stibine  is  not  evolved  by  the  action  of  strong-  KOH  upon  zinc  or  aluminum, 
nor  by  sodium  amalgam  in  neutral  or  alkaline  solution  (distinction  from  triad 
arsenic);  the  antimony  is  precipitated  as  the  metal  (Fleitmann,  -/.  C.,  18! 
329).  Stibine  is  slowly  oxidized  by  sulphur  to  Sb2S3  in  the  sunlight  at  ordinary 
temperature  and  rapidlj'  when  the  sulphur  (in  a  U  tube  mixed  with  glass  wool) 
is  heated  to  100°.  The  reaction  takes  place  according  to  the  following  equation: 
2SbH3  +  3S2  =  Sb,S3  +  3H2S  (Jones,  J.  C.,  1876,  29,  645). 

7.  Ignition. — By  ignition  in  the  absence  of  reducing  agents,  antimonic  acid 
and  anhydride  are  reduced  to  antimonous  antimonate.  Sb,O3.Sb.O5  or  Sb,O4 
(Sb'"SbvOJ,  a  compound  unchanged  at  a  dull  red  heat,  but  when  heated  to 
800°  this  oxide  is  further  reduced  to  antimonous  oxide  (ib). 

The  antimonates  of  the  fixed  alkali  metals  are  noi  vaporized  or  decomposed 
when  ignited  in  the  absence  of  reducing  agents;  hence,  by  fusion  in  the  crucible 
"with  sodium  carbonate  and  oxidizing  agents,  i.  e.,  with  sodium  nitrate  and  car- 
bonate, the  compounds  of  antimony  are  converted  into  non-volatile  sodium 
pyroantimonate,  Na4Sb2O7  ,  and  arsenic  compounds  if  present  are  at  the  same 
time  changed  to  sodium  orthoarsenate,  Na3AsO4  .  If  now  the  fused  mass  be 
digested  and  disintegrated  in  cold  water  and  filtered,  the  antimonate  is  sepa- 
rated as  a  residue,  Na;,H2Sb,O7  (4c),  while  the  arsenate  remains  in  solution 
with  the  excess  of  alkali.  The  operation  is  much  more  satisfactory  when  the 
arsenic  and  antimony  are  previously  fully  oxidized — as  by  digestion  with  nitric 
acid — as  the  oxidation  by  fusion  in  the  crucible  is  not  effected  soon  enough  to 
retain  all  the  arsenic  or  antimony  which  may  be  in  the  state  of  lower  oxides, 
sulphides,  etc.  If  compounds  of  tin  are  present  in  the  operation — and  it'  the 
fusion  is  not  done  with  excess  of  heat,  so  as  to  convert  sodium  nitrite  to  caustic 
soda  and  form  the  soluble  sodium  stannate — the  tin  will  be  left  as  stannic  oxide, 
SnO,  ,  in  the  residue  with  the  Na^H-Sb.O,  .  But  if  sodium  hydroxide  is  added 
in  the  operation,  the  tin  is  separated  as  stannate  in  solution  with  the  arsenic 
(Meyer,  ./.  C..  1849,  1,  388). 

All  compounds  of  antimony  are  completely  reduced  in  the  dry  way  on  char- 
coal with  sodium  carbonate,  more  rapidly  with  potassium  cyanide;  the  metal 
fusing  to  a  brittle  globule.  The  reduced  metal  rapidly  oxidizes,  the  white 
antimonous  oxide  rising  in  fumes,  and  making  a  crystalline  deposit  on  the 
support.  If  now  ammonium  sulphide  be  added  to  this  white  sublimate,  an 
orange  precipitate  is  a  sure  indication  of  the  presence  of  antimony  (Johnstone, 
C.  y.,  is$3.  58,  296).  The  same  white  oxide  is  formed  on  heating  antimony  or 
its  sulphides  in  a  glass  tube,  through  which  air  is  allowed  to  pass. 

8.  Detection. — Antimony  is  precipitated,  from  the  solution  acidulated 
with  hydrochloric  acid,  in  the  second  group  by  hydrosulphuric  acid  as  the 
sulphide  (6e).  By  its  solution  in  yellow  ammoniuin  sulphide  *  it  i?  sepa- 
rated from  Hg ,  Pb  ,  Bi ,  Cu  ,  and  Cd  .  In  the  Marsh  apparatus  the  anti- 
mony is  precipitated  on  the  Zn  as  the  metal,  a  portion  being  still  further 
reduced  to  stibine.  By  passing  the  gases,  stibine  and  arsine,  into  AgNO, 
solution,  the  antimony  is  precipitated  as  SbAg,  ,  antimony  argent  ide,  sepa- 

*  Antimony  as  sulphide  solution  in  potassium  sulphide  may  be  detected  electrolytically,  being 
deposited  as  Sb°.  Delicate  to  one  part  in  1,500,000  (Kohn,  J.  Soc.  Iwl.,  1891.  1 0,  327). 


§70,  10.  ANTIMONY.  31 

rating  it  from  the  arsenic  which  is  oxidized  and  passes  into  solution  as 
arsenous  acid.  The  SbAg3  is  dissolved  in  HC1  and  the  presence  of  the 
antimony  is  confirmed  by  the  precipitation  of  the  orange  colored  sulphide 
with  H2S .  Study  text  at  6  and  §84  to  §89.  For  distinction  between  Sbv 
and  Sb'"  see  §89,  7. 

9.  Estimation.— (1)  Tartaric  acid  and  water  are  added  to  SbCl3,  which  is 
then   precipitated   by   H2S   as   Sb2S   ,   and   after   washing   on   a  weighed   Gooch 
filter,  it  is  heated  to  230°  in  a  stream  of  CO2  ,  in  order  to  exclude  oxygen,  and 
weighed.      (2)   Antimonous  oxide,   sulphide,   or  any  oxysalt  of  antimony  is  first 
boiled  with  fuming  nitric  acid,  which  converts  it  into  SbjOs ,   and  then  by  ignition 
it  is  reduced  to  Sb2O4  ,  and  weighed  as  such.     (3)  The  trichloride  is  precipitated 
by  gallic  acid,   and  weighed  after  drying  at   100°.     (4)  In  the  presence  of  tin 
and  lead  oxidize  the  hydrochloric  acid  solution  of  the  salts  with  KC1O    (the  tin 
must  be  present  as  SnIV)  and  distil  in  a  current  of  HC1 .     The  stannic  and  anti- 
mony chlorides  are  volatile  (separation  from  lead).     To  the  distillate  add  metallic 
iron,  obtaining  stannous  chloride  and  metallic  antimony;    filter  and  wash  (sep- 
aration from  tin).     Fuse  the  precipitate  with  sodium  nitrate  and  sodium  car- 
bonate, digest  the  fused  mass  with  cold  water,  filter,  wash,  dry  and  weigh  as 
Na.:H2Sb2O7  (7)  (Tookey,  J.  C.,  1862,  15,  462;     and  Thiele,  A.,  1894,  263,  361). 

(5)  For  estimation  of  antimony  and  separation  from  arsenic  and  tin  by  the  use 
of  oxalic  acid,  see  Lessen  (Z.,  1888,  27,  218)  and  Clarke  (C.  N.,  1870,  21,  124). 

(6)  Volumetrically.     The    antimony    compound    is    converted    into    stibine    (6j) 
and  the  gas  passed  into  standard  silver  nitrate  solution.     The  solution  is  filtered 
and  the  excess  of  silver  nitrate  is  titrated  with  standard  sodium  chloride.     If 
arsenic  be  present  it  must  also  be  estimated  (§69,  9  (15)),  and  the  true  amount 
of  antimony  present  computed    from  the  two  determinations   (Houzeau,  J.  C., 
1873,  26,  407).     (7)  Sb'"  is  oxidized  to  Sbv  in  presence  of  NaHCO3  by  a  standard 
solution  of  iodine.     The  end  of  the  reaction  is  shown  by  the  blue  color  given  to 
starch.     (8)  Sb"'  is  oxidized  by  KC1O3  in  strong  HC1  solution  to  SbCls .     KI 
is  added,  which  reduces  the  Sb?  to  Sb'"  with  the  liberation  of  I2,  which  is  titrated 
with  Na-jS-jOs  solution.     (9)  Sb"'  is  oxidized  to   Sbv    in   presence   of  H2C4H4Oe 
by  KMnO4 .      (10)  Sb'"  is  oxidized  to  Sbv  by  K2Cr»O7 ,  and  the  excess  of  K2Cr2O7 
used  is  determined  by  a  standard  solution  of  FeSO4  ,  K3Fe(CN)6  being   used  to 
show  the  end  of  the  reaction.     (11)  The  antimony  as  the  triad  salt  is  treated  with 
an  excess  of  standard  K3Fe(CN)6  ;    the  excess  of  which  is  estimated  in  a  gas 
apparatus  with  H202  (Baumann,  Z,  angew,,  1892,  117), 

10.  Oxidation. — Stibine,  SbH3 ,  is  decomposed  by  heat  alone  into  anti- 
mony and  hydrogen  (6;).     By  burning  in  the  air  it  is  oxidized  to  Sb.,0; 
and  H20  .     Passed  into  a  solution  of  silver  nitrate,  SbAg3  is  produced,  or 
passed  into  a  solution  of  antimonous  chloride  or  potassium  hydroxide, 
sp.  gr.  1.25,  metallic  antimony  is  produced.     Excess  of  chlorine,  bromine, 
or  nitric  acid  in  presence  of  water  oxidizes  it  to  Sbv;  but  if  the  SbH3  be  in 
excess  metallic  antimony  is  precipitated.     With  excess  of  iodine  in  pres- 
ence of  water  Sb'"  is  produced;  if  the  stibine  be  in  excess  metallic  anti- 
mony.    Metallic  antimony  is  oxidized  by  nitric  acid,  chlorine  or  bromine 
to  Sb"'  or  Sbv,  depending  upon  the  amount  of  these  reagents  and  the 
temperature.     Iodine  oxidizes  the  metal  to  Sb'"  only,  except  in  alkaline 
mixtures  when  Sbv  is  formed. 

Antimonous  compounds  are  oxidized  to  antimonic  compounds  by  Cl , 
Br ,  HN03 ,  K2Cr207 ,  and  KMn04  ;  by  silver  oxide  in  presence  of  the  fixed 
alkalis  (6t);  by  gold  chloride  in  hydrochloric  acid  solution,  gold  being 


82  TIN.  §71,  1. 

deposited  as  a  yellow  precipitate  (§73,  10).  The  antimony  is  precipitated 
as  Sb205  unless  sufficient'  acid  be  present  to  dissolve  the  oxide:  4AuCl3  -f- 
3Sb203  -f  6H20  =  4Au  +  3Sb2Og  +  12HC1 . 

Antimonic  compounds  are  reduced  to  antimonous  compounds  by  HI  (6/) 
and  by  SnCL  (§69  and  §71,  10);  the  antimony  not  being  further  recuci'd 
(distinction  from  As).  Antimonic  and  antimonous  compounds  are  reduced 
to  the  metallic  state  by  Pb ,  Sn ,  Bi ,  Cu ,  Cd ,  Fe ,  Zn ,  and  Mg  ;  but  in 
the  presence  of  dilute  acids  and  metals  which  evolve  hydrogen  the  antimony 
is  still  further  reduced  to  stibine.  Iron  alone  or  in  the  presence  of  plat- 
inum (iron  platinum  wire  couple)  precipitates  the  antimony  from  acid  solu- 
tions as  Sb°  ;  0.000012  grams  can  be  detected  (Rideal,  C.  N.,  1885,  51,  292). 

Sodium  amalgam  with  dilute  sulphuric  acid  evolves  stibine  from  all 
antimony  solutions  (Van  Bylert,  B.,  1890,  23,  2968)  but  the  generation 
of  hydrogen  in  alkaline  solution,  i.  e.,  Zn  -f-  KOH ,  causes  the  reduction 
of  the  antimony  salt  to  the  metal  only,  in  no  case  evolving  stibine. 

§71.  Tin  (Stannum).    Sn  =  118.7.    Valence  two  and  four. 

1.  Properties. — Specific  gravity,  7.2984  (Rammelsberg,  B.,  1870,  3,  724);  welting 
point,  231.68°   (Callendar    and    Griffiths,    C.  N.,   1891,  63,  2).     Boils    at    2275° 
(Greenwood,  Proc.  Roy.  Soc.,  82,  396,   1908).     Does  not  distil  in  a  vacui m  at 
a  red  heat  (Schuller,  J.,  1884,  1550).     Tin  is  a  silver  white  metal,  does  not  t;  rnish 
readily   in   pure   air.     At   a   red   heat   it   decomposes   steam    with   evolution   of 
hydrogen;   at  a  white  heat  it  burns  in  the  air  with  a  dazzling  white  light,  fo:ming 
ShOo  .     It  is  softer  than  gold  and  harder  than  lead,  can  readily  be  hammered 
or  rolled  into  thin  sheets   (tinfoil);    at   100°  it   can  be  drawn  into  wire  a  id  at 
200°  can  be  pulverized.     Tin  possesses  a  strong  tendency  to  crystalline  struc- 
ture, and  when  bar  or  block  tin  is  bent  a  marked  decrepitation  "Zinngesc  hrei" 
(Levol,  A.  Ch.,  1859,  (3),  66,  110)  is  noticed,  due  to  the  friction  of  the  en  stals. 
Block  tin  exposed  to  severe  cold  (winter  of  1867-68,  at  St.  Petersburg,    --39°) 
crumbles  to  a  grayish  powder  (Fritsche,  B.,  1869,  2,  112),  considered  to  be  an 
allotropic  modification.     This  same  property  of  crumbling  is  noticed  in  samples 
of  tin  that  have  been  preserved  several  hundred  years   (Schertel,  J.  pr.,   1879, 
2,  19,  322).     The  grayish  powder  is  an  allotropic  modification  of  tin,  the  tran- 
sition temperature  being  20°  C.     Tin  forms  alloys  with  many  metals.     Eronze 
consists  of  copper  and  tin,  brass  frequently  contains  from  two  to  five  per  cent  of 
tin,  solder  consists  of  lead  and  tin.     All  the  easily  fusible  metals  as  Wood's  metal, 
etc.,  contain  tin.     For  many  references  concerning  tin  alloys,  see  Watts  (IV,  720). 

2.  Occurrence. — The  chief  ore  of  tin  is  cassiterite  or  tinstone,  a  nearly  pure 
crystallized    dioxide,    SnO2  ;     found    in    England,    Australia,    Malay    Peninsula, 
Bolivia,  Mexico,  and  to  a  very  limited  extent  in  the  United  States;    (Z).,  2,  1,  643). 
Stannite,  Cu2FeSnS4  ,  is  found  in  small  quantities  in  various  tin  veins. 

3.  Preparation. — The  reducing-  agent  employed  is  carbon.     The  impure  ore, 
SnO,  ,  is  first  roasted,  which  removes  some  of  the  arsenic  as  As2O3  ,  and  some 
of  the  sulphur  as  SO,  .     Then,  by  washing-,  the  soluble  and   some  of  t-.eju- 
soluble  impurities  are  washed  away,  the  heavier  SnO2  remaining-.     It  is  then 
fused  with  powdered  coal,  lime  being1  introduced  to  form  a  fusible  slag  with 
the  earthy  impurities.     It  is  refined  by  repeated  fusion.     Strictly  pure   ,in  is 
best  made  by  treating-  the  refined  tin  with  HN03  ,  and  then  reducing-  the  oxide 
thus  formed  by  fusion  with  charcoal;  or  by  reducing-  the  purified  chloride 

4.  Oxides  and  Hydroxides.— Tin  forms 'two  stable  oxides  and  cor  respo  id  ing- 
classes  of  salts:  starmous  oxide,  SnO,  black  or  blue  black,  and  stannic  c  xide. 
SnO2  ,  white;  the  latter  acts  both  as  a  base,  in  stannic  salts,  and  as  an  Anhy- 
dride, in  stannates.     Stan-nous  o.ridc  is  formed  (1)  by  precipitating  SnCl2  with 
K2CO8  ,  washing-  with  boiled  water  in  absence  of  air,  drying-  at  80°  or  1  >wer; 
then  dehydrating  by  heating-  in  an  atmosphere  of  hydrogen  or  carbon  dioxide 


§71,  o&.  TIN.  83 

(Loni?e,  r.  r.,  jssci,  ::i);  (2)  by  melting-  a  mixture  of  SnCl2  and  Na2C03  with 
.stirri.ig-  until  it  becomes  black,  and  removing  the  NaCl  by  washing  (Sandal, 
rhil.  May.,  1838,  (3),  12,  216;  Bottger,  /!.,  18.39,  29,  87).  titan-nous  hi/dro.ridc, 
Sn(C  I)  2*,  white  to  yellowish  white,  is  formed  by  adding-  alkalis  or  alkali 
carbc  nites  to  stannous  chloride,  washing  and  drying  at  a  low  temperature 
(Ditt  ;,  A.  Ch.,  1882,  (5),  27,  145).  (§12.) 

&(<•  iiic  oxide  exists  in  two  forms,  crystalline  and  amorphous.  The  native 
tinst  ne  is  nearly  pure  crystalline  SnO2  .  For  preparation  see  Bourgeois  (C.  r., 
1.887,  -04,  231)  and  Levy  and  Bourgeois  (('*.  r.,  1882,  94,  1305).  Amorphous  SnO2 
is  fo  tned  (1)  by  heating  tin  in  the  air  to  a  white  heat;  (2)  stannic  salts  are 
preci  'Hated  by  alkali  carbonates,  the  precipitate  washed  and  ignited;  (3)  tin 
is  DJ  dized  by  nitric  acid;  (4)  tin  filings  are  ignited  in  a  retort  with  HgO 
(/>.,  -  ,  1,  647).  Stannic  lii/dro.i-idc  or  stannic  acid  exists  in  two  forms:  (1)  Nor- 
mal i  "annic  acid,  SnO(OH)2  =  H2SnO:{  ,  is  formed  when  a  solution  of  stannic 
chloi  de  is  precipitated  by  barium  or  calcium  carbonate  (Freing,  Poyg.,  1842,  55, 
519);  if  an  alkali  carbonate  be  used  some  alkali  stannate  is  also  formed.  (2) 
Meta  tannic  acid,  H^Sn-.Oj,.  ,  is  formed  by  decomposition  of  tin  with  nitric 
acid  Hay,  C.  N.,  1870,  22,  298;  Scott,  C.  N.,  1870,  22,  322);  insoluble  in  acids  but 
chan  ed  on  standing  with  acids  to  normal  stannic  acid,  which  is  readily  soluble 
in  a<  ds  (56).  It  is  also  formed  when  stannic  chloride  is  boiled  in  concen- 
trate I  solution  with  most  of  the  alkali  salts:  5SnCl4  +  20Na2S04  -4-  15H20  = 
H10S.i5016  +  20NaCl  +  2()NaHS04  ,  or  according  to  Fresenius  (16th  edition), 
271:  SnCl4  +  4Na2SO,  +  4H2O  =  Sn(OH)4  +  4NaCl  +  4NaHSO4 .  It  is  also 
form-  d  together  with  hydrochloric  acid  when  stannic  chloride  is  boiled  with  a 
large  excess  of  water. 

5.  Solubilities. — a. — Metal. — Tin  dissolves  in  hydrochloric  acid  slowly  when 
the  i  "Ad  is  dilute  and  cold,  but  rapidly  when  hot  and  concentrated,  stannous 
chlor  de  and  hydrogen  being  produced  (a);  in  dilute  sulphuric  acid  slowly,  with 
separation  of  hydrogen  (6),  (not  at  all  even  in  hot  acid  if  more  dilute  than 
H2SC1.6H2p  (Ditte,  A.  Ch.,  (5),  27,  145);  in  hot  concentrated  sulphuric  acid 
rapid  y,  with  separation  of  sulphurous  anhydride  and  sulphur  (c);  nitric  acid, 
rapid,  y  converts  it  into  metastannic  acid,  insoluble  in  acids  (d)'}  very  dilute 
nitric  acid  dissolves  it  without  evolution  of  gas  as  stannous  nitrate  and  ammo- 
nium nitrate  (e]  (Maumene,  BL,  (2),  36,  598);  nitro-hydrochloric  acid  dis- 
solves tin  easily  as  stannic  chloride  (/),  potassium  hydroxide  solution  dissolves 
it  ve.:y  slowly,  and  by  atmospheric  oxidation  (p);  or,  at  high  temperatures, 
with  evolution  of  hydrogen  (h).  Bromine  vapors  readily  attack  melted  tin 
with  formation  of  SnBr<  ,  colorless  crystals,  melting  point  30°  (Carnelley  and 
O'Shca,  J.  C.,  1878,  33,  55).  Dry  chlorine  gas  attacks  tin  readily  in  the  cold, 
producing  stannic  chloride  as  vapor  or  colorless  liquid.  The  action  is  vigorous 
enough  in  strong  chlorine  to  produce  a  flame. 

Sn  +  2HC1  =  SnCL  +  H2 

Sn  +  H2S04  —  SnS04  +  H2 

Sn  +  2H2S04  =  SnSO,  +  2H2O  +  S02 
and  then  4SriS04  +  2SO,  +  4H2S04  =  4Sn(S04)2  +  S2  +  4H2O 

(d)  15Sn  +  20HN03  +  5H20  =  3H10Sn5O15  +  20NO 

(e)  4Sn  +  10HN03  =  4Sn(NO3)2  +  3H.O  +  NH4N08 
(0      Sn  +  2C12  =  SnClt 

(g)     2Sn  +  4KOH  +  02  =  2K2SnO2  +  2H2O 
(7t)     Sn  -f  2KOH  =  K2Sn02  +  H2 

6. — Oxides. — Stannous  oxide  is  insoluble  in  water,  soluble  in  acids  (Ditte,  A.  Ch., 
1882,  (5),  27,  145;  Weber,  /.  C.,  1882,  42,  1266),  oxidized  by  nitric  acid  when 
heat?d,  forming  the  insoluble  metastannic  acid.  Sidiinous  hydroxide  is  readily 
soluble  in  all  the  solvents  of  the  oxide,  and  is  also  readily  soluble  in  fixed 
alkali  hydroxides.  Sia-nnic  o.i-ide,  Sn02  ,  is  insoluble  in  water;  soluble  with 
difficulty  in  alkalis;  insoluble  in  acids  except  in  concentrated  H2SO4  (D.,  2,  1, 
648).  Sulphur  forms  SnS2  and  SO-,;  chlorine  forms  SnCl4  (Weber,  Pogg.,  18(51, 
112,  619).  Normal  stinuiic  acid,  BLSnOj  ,  freshly  precipitated,  is  soluble  in 

•According  to  other  authorities  Sn(OH)8  does  not  exist,  hut  a  liydratccl  oxide  is  formed, 
SnO .  Sii(OH)a  (Graham-Otto,  a,  2, 1207 ;  D.,  2, 1,  657;  Graelin-Kraut,  «,  107). 


84  TIN.  §71,  5c. 

fixed  alkali  hydroxides  and  in  acids  (Ditte,  C.  r.,  1887,104,  172);  insoluble  in 
water  and  changed  by  hot  nitric  acid  to  the  insoluble  metastannic  acid.  Mela- 
stannic  acid,  H10Sn5Oio,  is  insoluble  in  water  and  acids,  HC1  changes  it  to 
metastannic  chloride  insoluble  in  the  acid,  but  soluble  in  water  after  removal 
of  the  acid;  soluble  in  the  fixed  alkalis  as  metastannates,  which  are  soluble  in 
water  and  precipitated  by  acids.  Metastannic  acid  in  contact  with  HC1  is 
gradually  changed  to  stannic  acid  (Barfoed,  J.  pr.,  1867,  101,  368). 

C. — Salts. — The  sulphides  and  phosphates  of  tin  are  insoluble  in  water,  also 
stannous  oxychloride;  stannous  sulphate,*  bromide  and  iodide;  and  stannic 
chloride  and  bromide  dissolve  in  pure  water  with  little  or  no  decomposition 
(Personne,  C.  r.,  1862,  54,  216;  and  Carnelley  and  O'Shea,  J.  C.,  1878,  33,  55). 
Stannous  chloride  is  soluble  in  less  than  two  parts  of  water  (Engel,  A.  Ch.,  1891, 
(6),  17,  347);  but  more  water  decomposes  it,  unless  a  strong  excess  of  acid  be 
present:  2SnCl2  -f-  H20  =  SnO.SnCl2  +  2HC1 .  Pure  stannic  chloride  is  a 
liquid;  sp.  gr.,  2.2;  boiling  point,  114°  (Walden,  Z.  ph.  Ch.,  43);  solidifies  at 
—  33°  (Besson,  C.  r.,  1889,  109,  940).  The  liquid  combines  with  water,  liberating 
heat  to  form  crystals  of  SnCl.).3H2O  ,  which  are  readily  soluble  in  excess  of  water 
(D.,  2,  1,  662).  Stannic  chloride  is  completely  decomposed  by  boiling  water. 
The  nitrates  of  tin  are  very  easily  decomposed  by  water  and  require  free  acid 
to  keep  them  in  solution  (Weber,  J.  pr.y  1882,  (2),  26,  121;  Montemartini, 
Gazzetta,  1892,  22,  384).  Stannic  iodide  is  readily  soluble  in  water  (Schneider, 
Pogg.,  1866,  127,  624).  Stannic  sulphate  is  easily  soluble  in  water,  but  is  de- 
composed by  a  large  excess  (Ditte,  C.  r.,  1887,  104,  171).  Stannous  and  stannic 
chloride,  and  stannic  iodide  are  soluble  in  alcohol.  Stannous  nitrate  and  stannic 
sulphate,  and  bromide  are  deliquescent.  Stannous  sulphide  is  insoluble  in  water, 
soluble  in  HC1  with  formation  of  H2S;  decomposed  by  HNO3  with  oxidation  to 
metastannic  acid;  insoluble  in  solution  of  the  normal  alkali  sulphides,  but  soluble  in 
the  polysulphides  with  oxidation  to  a  stannic  compound  (6e).  Stannic  sulphide  is 
soluble  in  HC1 ,  with  evolution  of  H2S  ;  and  in  solutions  of  the  alkali  sulphides. 

6.  Reaction's. — a.  Alkali  hydroxides  and  carbonates  precipitate  from  solu- 
tions of  stannous  salts,  stannous  hydroxide,  Sn(OH)2  (4) ,  white,  readily  soluble 
in  excess  of  the  fixed  alkali  hydroxides,  insoluble  in  water,  ammonium  hy- 
droxide and  the  alkali  carbonates  (distinction  from  antimony).  It  is  also 
precipitated  by  barium  carbonate  in  the  cold  (Schaffner,  A.,  1844,  51,  174). 
SnCl2  +  2KOH  —  Sn(OH)2  +  2KC1 

Sn(OH)2  +  2KOH  =  K2Sn02  +  2H20 
SnCl2  +  4KOH  —  K2SnO,  +  2KC1  +  2H20 
SnCl2  +  Na2COs  +  H2O  =  Sn(OH)2  +  2NaCl  +  C02 

By  gently  heating  the  solution  of  potassium  stannite,  K2Sn02 ,  crystalline 
stannous  oxide,  SnO ,.  is  formed.  By  rapid  boiling  of  a  strong  potassium 
hydroxide  solution  of  stannous  hydroxide  part  of  the  tin  is  oxidized  and 
the  remainder  precipitated  as  metallic  tin;  2K2Sn02  +  ^20  =  Sn  -f- 
K2Sn03  +  2KOH .  The  reaction  proceeds  more  rapidly  upon  the  addition 
of  a  little  tartaric  acid.  Stannic  salts  are  precipitated  by  alkali  hydroxides 
and  carbonates  as  stannic  acid,  H2Sn03  soluble  in  excess  of  the  fixed  alkali 
hydroxides,  insoluble  in  ammonium  hydroxide  and  the  alkali  carbonates 
(Ditte,  A.  Ch.,  1897  (6),  30,  282). 

SnCl<  +  4KOH  =  H2SnO8  +  4KC1  +  H20 

H2Sn08  +  2KOH  =  K2SnO,  +  2H2O 
SnCl4  +  6KOH  =  K2Sn08  +  4KC1  +  3H2O 
SnCl,  +  2Na2C08  +  H20  =  H2Sn08  +  4NaCl  +  2CO2 

*  Stannous  sulphate  is  decomposed  by  an  excess  of  cold  water  forming  2Sn8O4.4SnO.8HaO| 
and  by  a  small  amount  of  hot  water  forming  SnSO^.gSnO  (Ditte,  A.  Ch.,  1883,  (5),  27, 161). 


|71,  6«.  TIN.  85 

Metastannic  salts  are  precipitated  as  metastannic  acid  soluble  in  potassium 
hydroxide  not  too  concentrated,  not  readily  soluble  in  sodium  hydroxide, 
insoluble  in  ammonium  hydroxide  excepting  when  freshly  precipitated  in 
the-  cold,  and  the  alkali  carbonates. 

6. — Oxalic  acid  forms  a  white  crystalline  precipitate  with  a  nearly  neutral 
solution  of  staiinous  chloride,  soluble  in  hydrochloric  acid,  not  readily  soluble 
in  ammonium  chloride.  If  a  nearly  neutral  solution  of  stannous  chloride  be 
added  drop  by  drop  to  a  solution  of  ammonium  oxalate,  the  white  precipitate 
which  forms  at  once  dissolves  in  the  excess  of  the  ammonium  oxalate.  Stannic 
chloride  is  not  precipitated  by  oxalic  acid  or  ammonium  oxalate  (Hausmann 
and  Loewenthal,  A.,  1854,  89,  104). 

Potassium  cyanide  precipitates  both  stannous  and  stannic  salts,  white,  in- 
soluble in  excess  of  the  cyanides.  Potassium  ferrocyanide  precipitates  from 
stannous  chloride  solution  stannous  ferrocyanide,  Sn2Fe(CN)6  ,  white,  insoluble 
in  water,  solublp  in  hot  concentrated  hydrochloric  acid.  From  stannic  chloride 
is  precipitated  a  greenish  white  gelatinous  precipitate,  soluble  in  hot  hydro- 
chloric acid,  but  reprecipitated  upon  cooling  (distinction  from  antimony) 
(Wyrouboff,  A.  Ch.,  1876,  (5),  8,  458).  Potassium  ferricyanide  precipitates  from 
solutions  of  stannous  chloride,  stannous  ferricyanide,  Sn3(Fe(CN)0)2  ,  white, 
readily  soluble  in  hydrochloric  acid.  On  warming,  the  ferricyanide  is  reduced 
to  ferrocyanide  with  oxidation  of  the  tin.  No  precipitate  is  formed  by  the 
ferricyanide  with  stannic  chloride. 

c. — The  nitrates  of  tin  are  not  stable.  Stannous  nitrate  is  deliquescent  and 
soon  decomposes  on  standing  exposed  to  the  air.  Stannous  salts  when  heated 
with  nitric  acid  are  precipitated  as  SnO3;  but  if  stannous  chloride  be  warmed 
with  a  mixture  of  equal  parts  of  nitric  and  hydrochloric  acids,  stannic  chloride 
and  ammonium  chloride  are  formed  (Kestner,  A.  Ch.t  1860,  (3),  58,  471). 

df. — Hypophosphorous  acid  does  not  form  a  precipitate  with  stannous  or 
stannic  chlorides,  nor  are  these  salts  reduced  when  boiled  with  the  acid.  Sodium 
hypophosphite  forms  a  white  precipitate  with  stannous  chloride,  soluble  in 
excess  of  hydrochloric  acid;  no  precipitate  is  formed  with  stannic  chloride. 
Phosphoric  acid  and  soluble  phosphates  precipitate  from  solutions  of  stannous 
salts,  not  too  strongly  acid,  stannous  phosphate,  white,  of  variable  composition, 
soluble  in  some  acids  and  KOH;  insoluble  in  water  (Lenssen,  A.,  1860,  114, 
113).  With  stannic  chloride  a  white  gelatinous  precipitate  is  formed,  soluble 
in  HC1  and  KOH  ,  insoluble  in  HNOS  and  HC,H802  .  If  the  stannic  chloride  be 
dissolved  in  excess  of  NaOH  before  the  addition  of  Na2HPO4  and  the  mixture 
then  acidulated  with  nitric  acid,  the  tin  is  completely  precipitated  as  stannic 
phosphate  (separation  from  antimony).  However,  the  precipitate  always  car- 
ries a  little  antimony  (Bornemann,  Z.  angew.,  1899,  635). 

e.  Hydrosulphuric  acid  and  soluble  sulphides  precipitate  from  solutions 
of  stannous  salts  dark  brown  hyd rated  stannous  sulphide,  SnS  (a),  insol- 
uble in  dilute,  soluble  in  moderately  concentrated  HC1  (b).  It  is  readily 
dissolved  with  oxidation  by  alkali  supersulphides,  the  yellow  sulphides, 
forming  thiostannates  (c);  from  which  acids  precipitate  the  yellow  stannic 
sulphide  (d).  The  normal,  colorless  alkali  sulphides  scarcely  dissolve  any 
stannous  sulphide  at  ordinary  temperature,  compare  (§69,  6e  and  §70,  6e), 
but  hot  concentrated  K2S  dissolves  SnS  forming  K2SnS3  and  Sn  (e)  (Ditte, 
C.  r.,  1882,  94,  1419;  Baubigny,  J.  C.,  1883,  44,  22).  Potassium  and 
sodium  hydroxides  dissolve  it  as  stannites  and  thiostannites  (f),  from 
which  acids  precipitate  again  the  brown  stannous  sulphide  (g).  Am- 
monium hydroxide  and  the  alkali  carbonates  do  not  dissolve  it  (distinction 
from  arsenic,  §69,  6e).  The  insolubility  in  fixed  alkali  carbonates  is  a 


86  2W.  §71, 6/. 

distinction  from  antimony  (§70,  6e).  Nitrohydrochloric  acid  (free  chlorine) 
dissolves  it  as  stannic  chloride,  with  residual  sulphur  (h).  Nitric  acid 
oxidizes  it  to  raetastannic  acid  without  solution  (t)  (separation  from 
arsenic,  §69,  6e). 

(a)     SnCl2  +  H2S  =  SnS  +  2HC1 

(&)     SnS  +  2HC1  =  SnCL  +  ELS 

(c)  SnS  +  (NH4)2S,  =  (NH.t)2"snS3 

(d)  (NH4)2SnS8  +  2HC1  =  SnS2  +  2NH4C1  +  H2S 

(e)  2SnS  +  K2S  =  K2SnS3  +  Sn 

(f)  2SnS  +  4KOH  =  K2Sn02  +  K2SnS2  +  2H2O 

(0)     (K,Sn02  +  K2SnS2)  +  4HC1  =  2SnS  +  4KC1  +  2H20 

(7i)     2SnS  +  4C12  =  2SnCl4  -f  S2 

(?)      SOSnS  +  40HN03  +  10H20  =  GH10Sn5015  +  40NO  +  15S2 

Solutions  of  stannic  salts  are  precipitated  as  stannic  sulphide,  SnS2 . 
hydrated,  yellow,  having  much  the  same  solubilities  as  those  given  for 
stannous  sulphide,  with  this  difference,  that  stannic  sulphide  is  moderately 
soluble  in  normal,  colorless,  alkali  sulphides.  The  following  equations 
illustrate  the  most  important  reactions: 

SnCl4  +  2H2S  =  SnS2  +  4HC1 

SnS2  +  4HC1  =  SnCl4  +  2H2S 

SnS2  +  (NH4)2S  =  (NH4)2SnS3 

2SnS2  +  2(NH4)2S2  =  2(NH4)2SnS3  +  S2 

3SnS2  +  6KOH  =  K2SnO3  +  2K2SnS3  +  3H2O 

(K2Sn03  +  2K2SnS3)  +  6HC1  =  3SnS3  +  GKC1  +  3H2O 

SnS2  +  2C12  =  SnCl4  +  S2 

15SnS2  +  20HN03  +  5H20  =  3H10Sn5015  +  15S2  +  20NO 

Sodium  thiosulphate  does  not  form  a  precipitate  with  the  chlorides  of  tin 
(separation  from  As  and  Sb)  (Lesser,  Z.,  1888,  27,  218).  Sulphurous  acid  and 
sodium  sulphite  precipitate  from  stannous  chloride  solution  not  too  strongly 
acid,  sta-nnous  sulphite,  SnS03  ,  white,  readily  soluble  in  HC1 .  When  wariaed  in 
the  presence  of  hydrochloric  acid,  si^lphur  dioxide  acts  as  an  oxidizing  agent 
upon  the  stannous  salt.  A  precipitate  of  Sn6O10S2  or  SnS2  is  formed,  or  H2S 
is  evolved  and  SnCl4  formed,  depending  upon  the  amount  of  HC1  present. 

6SnCl2  +  2S02  +  GH20  =  Sn6O10S2  +  12HC1 
6SnCl2  +  2S02  +  8HC1  =  SnS2  +  5SnCl4  +  4H2O 
3SnCl2  +  S02  +  6HC1  =  3SnCl4  +  H2S  +  2H2O 

Stannic  chloride  does  not  give  a  precipitate  with  sulphurous  acid  or  sodium 
sulphite. 

The  sulphates  of  tin  are  formed  by  dissolving  the  freshly  precipitated 
hydroxides  in  sulphuric  acid  and  evaporating  at  a  gentle  heat.  They  cannot  be 
formed  by  precipitation  and  are  decomposed  by  water  (Ditte,  A.  CJi.,  1882,  (5), 
27,  145). 

f. — Potassium  iodide  added  to  a  concentrated  water  solution  of  stannous  chlo- 
ride forms  first  a  yellow  precipitate  soluble  in  excess  of  the  SnCl,  .  Further 
addition  of  KI  gives  a  yellow  precipitate  rapidly  turning  to  dark  orange  needle- 
like  crystals,  often  forming  in  rosette-like  clusters.  If  a  drop  of  the  stannous 
chloride  solution  be  added  to  an  excess  of  potassium  iodide  the  yellow  p^ecipi- 
tate  is  formed,  which  remains  permanent  unless  a  further  quantity  of  sts  nnous 
chloride  be  added  when  the  orange  precipitate  is  formed.  The  orange  p  -ecipi- 
tate  is  probably  SnI,  ,  and  is  soluble  in  HC1  ,  KOH  ,  and  C,H5OH  ,  soluble  in 
large  excess  of  KI  and  sparingly  soluble  in  H2O  with  some  decomposition. 


§71,  7.  TIN.  87 

The  yellow  precipitate  is  probably  a  double  salt  of  stannous  iodide  and  potas- 
sium iodide,  and  has  about  the  same  solubilities  as  the  orange  precipitate 
(Personue,  J.,  1862,  171;  Boullay,  A.  (7/,,  1K27,  (2),  34,  372).  Potassium  iodide  in 
concentrated  solution  precipitates  xtumnc  iodide,  yellow,  from  very  concentrated 
\vate>  solutions  of  stannic  chloride.  The  precipitate  is  readily  soluble  in  water 
to  a  i  olorless  solution  (Schneider,  </.,  I860,  229).  Hydriodic  acid  does  not  give 
free  7.  with  Sniv  ,  distinction  from  Sbv  and  Asv  (Harroun,  J.  C.,  1882,  42,  GG1). 

The  chlorates,  bromates  and  iodates  of  tin  have  not  been  thoroughly  studied 
(Watti,  1,  539,  III.,  22;  D.,  2,  1,  675).  Stannous  chlorate  appears  to  be  formed 
when  potassium  chlorate  is  added  to  a  concentrated  water  solution  of  stannous 
chlor.de;  it  dissolves  on  addition  of  HC1,  and  nearly  all  dissolves  in  excess  of 
water.  With  KBrO3  ,  bromine  is  liberated,  and  with  KI03  iodine  is  liberated. 
Potassium  chlorate,  bromate  and  iodate  all  form  precipitates  with  stannic 
chlor  de,  soluble  in  HC1  without  liberation  of  the  halogen. 

g.—  Stannous  arsenate,  2SnO.As205  ,  a  voluminous  flocculent  precipitate  is 
forme  d  by  adding  a  solution  of  SnCl2  to  a  concentrated  acetic  acid  solution  of 
K3AsO4  ,  decomposed  by  heating  to  As  ,  As203  and  Sn02  (Lenssen,  A.,  I860,  114, 
115).  Stannic  arsenate,  2SnO,.As2O5  ,  a  white  gelatinous  precipitate  is  formed 
by  acVjling  HNO3  to  a  mixture  of  Na2Sn03  and  Na3AsO4  (Haeffely,  J.,  1855,  395). 
With  antimony,  tin  acts  as  a  base,  forming  stannous  and  stannic  antimonites 
and  antimonates  (Lenssen,  I.  c.). 

h. — If  potassium  chromate  be  dropped  into  a  hydrochloric  acid  solution  of 
stannous  chloride  there  is  immediate  reduction  of  chromium  with  formation 
of  a  dirty  brown  precipitate.  If  stannous  chloride  be  carefully  added  to  potas- 
sium chromate  in  excess,  an  abundant  yellowish  precipitate  is  obtained  without 
much  apparent  reduction  of  the  chromium.  Potassium  chromate  added  to 
stannic  chloride  gives  an  abundance  of  bright  yellow  precipitate  soluble  in 
excess  of  SnCl4  ,  insoluble  in  H2O  ,  soluble  with  difficulty  in  HC1 .  K2Cr2O7 
also  gives  a  precipitate  with  SnCl2  and  SnClt  (Leykauf,  J.  pr.,  1840,  19,  127). 

i.  An  ammoniacal  solution  of  silver  nitrate  is  reduced  to  metallic  silver 
by  a  solution  of  potassium  stannite.  The  reagent  (silver  nitrate  solution 
one  part,  to  ammonium  hydroxide  sixteen  parts)  serves  as  a  delicate  test 
for  the  presence  of  Sn"  in  solution  in  KOH .  The  addition  of  KOH  in 
excess  to  an  unknown  solution  removes  all  heavy  metals  except  Pb ,  Sb  , 
Sn ,  Al ,  Cr ,  and  Zn  ;  of  these  tin  only  precipitates  metallic  silver  from  the 
strongly  ammoniacal  solution  in  the  cold.  Antimonous  and  arsenous 
compounds  give  the  black  precipitate  of  metallic  silver  if  the  solution  be 
boileJ. 

;.  A  solution  of  mercuric  chloride,  HgCl2 ,  reacts  with  stannous 
chloride  solution,  forming  SnCl4  and  a  precipitate  of  HgCl  (white)  or  Hg°, 
gray,  depending  upon  the  relative  amounts  present  (§58,  6g). 

k.  Stannous  salts  react  with  (NH4)2Mo04 ,  giving  a  blue-colored 
solution  of  the  lower  oxides  of  molybdenum,  constituting  a  delicate  test 
for  Sn"  (§75,  60). 

7.  Ignition. — Before  the  blow-pipe,  on  charcoal,  with  sodium  carbonate,  and 
more  readily  by  addition  of  potassium  cyanide,  tin  is  reduced  to  malleable 
lustrous  globules — brought  to  view  (if  minute,  under  a  magnifier)  by  repeated 
trituration  of  the  mass  with  water,  and  decantation  of  the  lighter  particles. 
A  little  of  the  white  incrustation  of  stannic  oxide  will  collect  on  the  charcoal 
near  the  mass,  and,  by  persistence  of  the  flame  on  the  globules,  the  same  coat- 
ing forms  upon  them.  This  coating,  or  oxide  of  tin,  moistened  with  solution  of 
cobalt  nitrate,  and  again  ignited  strongly,  becomes  of  a  blue-green  color.  SnOa 
fused  with  KCN  gives  metallic  tin  (Bloxam,  J,  C.,  1865,  18,  97), 


88  TIN.  §71,  8. 

8.  Detection. — Tin  is  precipitated,  from  the  solution  acidulated  with 
hydrochloric  acid,  in  the  second  group  by  hydrosulphuric  acid,  as  the  sul- 
phide (6e).     By  its  solution  in  yellow  ammonium  sulphide  it  is  separated 
from  the  Copper  Group  (Hg,  Pb ,  Bi,  Cu,  and  Cd).     By  the  reaction  in 
the  Marsh  apparatus  the  tin  is  reduced  to  the  metal  and  is  not  dissolved 
as  long  as  zinc  is  still  present.     The  residue  Sn  (Zn ,  Sb ,  Au ,  and  Pt)  in 
the  Marsh  apparatus  is  warmed  with  hydrochloric  acid,  which  dissolves 
the  Sn  as  SnCl2 .     This  is  detected  by  its  reducing  action  on  HgCl2 ,  giving 
a  white  precipitate  of  HgCl  or  a  gray  one  of  Hg°  (6/). 

A  short  test  for  the  detection  of  tin  in  the  stannous  condition,  or  after 
its  reduction  to  that  condition,  consists  in  treating  the  solution  with  an 
excess  of  cold  KOH  (separation  of  Pb ,  Sn ,  Sb  ,  Al ,  Cr,  and  Zn,  from 
all  other  heavy  metals) ;  and  adding  to  this  solution,  filtered  if  necessary, 
a  solution  of  AgN03  in  a  great  excess  of  NH4OH  (one  part  AgN03  to  sixteen 
parts  NH4OH).  A  brown -black  precipitate  of  metallic  silver  indicates 
that  tin  was  present  in  the  stannous  condition  (6t).  Consult  also  §90 
and  §92. 

9.  Estimation. — (1)  Gravimetrically.    It  is  converted  into  SnO2  ,  and   after 
ignition  weighed.     (2)  Volumetrically.     To  SnCl2  add  KNaC4H4Os  and  NaHCO, , 
then  some  starch  solution  and  a  graduated  solution  of  iodine,  until  a  perma- 
nent blue  coloration  appears.     (3)   To  SnCl2   add  slight  excess  of  FeCl,  ,  and 
determine  the  amount  of  FeCl2  formed,  by  a  graduated  solution  of  KMn04  . 
(4)  By  electrolytic  deposition  from  a  solution  of  the  double  oxalate,  rendered 
slightly  acid  with  oxalic  acid. 

10.  Oxidation. — Metallic  tin  reduces  solutions  of  Ag ,  Hg ,  Bi ,  Cu  ,  Pt , 
and  Au ,  to  the  metallic  state.  Sn"  is  oxidized  to  SnIV  by  free  HN02 , 
HNOg1,  H3Fe(CN)6 ,  H2S03  and  H2S04  (if  hot),  Cl ,  HC10 ,  HC102 ,  HC10, , 
Br ,  HBr03 ,  1 3,  and  HI03 .  Also  by  Pb"  (in  alkaline  solution  only),  PbIV  , 
Ag'2,  Hg',  Hg",  Asv,  As'"  (in  presence  of  HC1),  Sbv,  Movl,  Bi"',  Cu', 
Pd(N03)2 ,  PtIV  4,  Fe'",  FeVI,  CrVI,  Co'",  Ni'",  and  Mn2+n.  Chlorine,  bromine 
and  iodine  act  more  vigorously  in  alkaline  than  in  acid  mixtures.  The 
above  mentioned  metallic  forms  oxidize  Sn"  in  both  acid  and  alkaline 
mixtures. 

Stannous  chloride,  is  one  of  the  most  convenient  and  efficient  of  the 
ordinary  discriminative  deoxidizing  agents  for  operations  in  the  wet  way. 
As  stannic  chloride  is  soluble  in  the  solvents  of  stannous  chloride  no 
precipitate  of  tin  is  made  by  its  reducing  action;  but  many  other  metals 
are  so  precipitated  by  reduction  to  insoluble  forms,  and  are  thus  identified 
in  analysis,  e.  g.,  mercuric  chloride  is  reduced  from  solution,  first  to  white 
mercurous  chloride,  and  then  to  gray  mercury  (detection  of  mercury); 
silver  nitrate,  to  brown-black  silver  (detection  of  tin);  all  soluble  com- 

» Kestner,  A.  C7i.,  I860,  (3),  58,  471.    'Ditte,  A.  Ch.,  1882,  (6),  27, 145.    «  Thomas,  C.  r.,  1896, 122, 
1639.    «  Ditte,  C.  r.,  1882,  94, 1114. 


§71,  10.  TIN.  89 

pounds  of  arsenic  in  strong  HC1  (detection  of  arsenic);  bismuth  salts,  to 
metallic  bismuth  (in  alkaline  mixture  §76,  6*7);  and  ferric  salts,  to 
ferrous  salts,  left  in  solution,  much  used  in  volumetric  analysis  of  iron 
(9,  and  §126,  Qg  and  9);  auric  chloride  is  reduced  to  the  metal  by  stannous 
chloride,  forming  a  colored  precipitate  varying  from  brown  to  reddish- 
brown  or  purple-red  according  to  the  amount  of  stannic  chloride  present. 
This  finely  divided  precipitate  of  gold  is  called  "  Purple  of  Cassius  "  (Max 
Muller,  J.  pr.,  1884,  30,  252). 

Solutions  of  SnIV  and  Sn"  are  reduced  to  the  metallic  state  by  Cd ,  Al . 
Zn ,  and  Mg .  According  to  Rideal  (C.  N.,  1885,  51,  292)  0.00003  grams 
of  tin  in  solution  may  be  detected  as  the  metal  by  reduction,  using  the 
gold  zinc  wire  couple.  Stannic  salts  are  reduced  to  stannous  salts  by 
metallic  tin,  copper  or  iron  (Allen,  J.  (7.,  1872,  25,  274). 


90 


REACTIONS  OF  ARSENIC,  ANTIMONY  AND   TIN. 


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§73,  5a.  GOLD.  91 

§73.  Gold   (Aurnm)    Au=  197.2.     Valence  one  and  three. 

1.  Properties.—  Specific  gravity,   19.30  to   19.34   (Rose,  Pogg.,   1848,  76,  403). 
Melting  point,  1063°  (Cr.  B.  8.,  35,  1915).     It  is  a  yellow  metal,  that  from  dif- 
ferent parts  of  the  world  varying  slightly  in  color;    the  presence  of  very  small 
traces  of  other  metals  also  affects  the  color.     It  is  softer  than  silver  and  harder 
than  tin;   possesses  but  little  elasticity  or  metallic  ring.     It  is  the  most  malleable 
and  ductile  of  all  metals;    one  gram  can  be  drawn  into  a  wire  2000  metres  long. 
The  presence  of  other  metals  diminishes  the  ductility.     It  may  be  rolled  into 
sheets  0.0001  mm.  thick.     At  a  very  high  heat  it  vaporizes  (DeviUe  and  Debray, 
A.  Ch  ,  1859,  (3),  56,  429).     It  is  a  good  conductor  of  electricity,  equal  to  copper, 
not  so  good  as  silver.     It  has  a  high   coefficient  of  expansion   and  cannot  be 
moulded  into  forms  but   must  be  stamped.     On   account   of  its  softness,   gold 
is  seldom   used   absolutely  pure,   but  is  hardened  by  being  alloyed   with  other 
incl.-ils,  as  Ag  ,  Cu  ,  etc. 

2.  Occurrence. — Gold  is  usually  found  native,  but  never  perfectly  pure,  being 
always  alloyed  with  silver,  and  occasionally  also  with  other  metals.       It  is  found 
as  gold-dust  in  alluvial  sand,  sometimes  in  nuggets,  and  sometimes  disseminated 
in  veins  of  quartz. 

3.  Preparation. — (1)  Washing.     This    consists   in  treating  the  well-powdered 
ore  with  a  stream  of  water,  the  heavy  gold  settling  to  the  bottom.     (2)  Amalga- 
mation.    This  consists  in  dissolving  the  gold  in  mercury  and  then  separating 
it  from  the  latter  by  distillation.     (3)  By  fusing-  with  metallic  lead,  which  dis- 
solves the  gold,  the  liquid  alloy  settling-  to  the  bottom  of  the  slag1.     The  gold  is 
afterward  separated  from  the  lend  by  cupellation.     The  silver  is  separated  from 
the  gold  by  dissolving'  it  in  nitric  or  sulphuric  acid.     Or  the  whole  is  dissolved 
in  nitrohydrochloric  acid,  and  the  gold  precipitated  in  the  metallic  state  by 
some    reducing-    agent;    ferrous   sulphate    being    usually    employed.      Another 
method  is  to  pass  chlorine  into  the  melted  alloy.     The  silver  chloride  rises  to 
the  surface,  while  the  chlorides  of  Zn  ,  Bi  ,  Sb  ,  and  As  (if  present)  are  vola- 
tilized, and  the  pure  gold  remains  beneath.     A  layer  of  fused  borax  upon  the 
surface  prevents  the  silver  chloride  from  volatilizing.     (4)  By  treatment  with  a 
solution  of  KCN  .     (£)  By  amalgamation  with  mercury  and  electrolysis  at  the 
same  time. 

4.  Oxides  and  Hydroxides. — Aurous  oxide,  Au2O  ,  is  very  unstable,  heating  to 
about  250°,  decomposing  it  into  the  metal  and  oxygen.     The  hydroxide  is  pre- 
pared by  reducing  the  double  bromide  with  SO2  in  ice-cold  solution;  heating  to 
200°,   changing  it  to  the  oxide   (Kriiss,   A.,   1886,  237,   274).     Auric  hydroxide, 
Au(OH).,  ,  is  prepared  by  precipitation  from  the  chloride  solution  with  MgO 
(Kriiss,  7.  e.).     It  is  a  yellow  to  brown  powder,  changing  to  the  oxide  upon  dry- 
ing at  100°.     Heating  to  230°  gives  the  metal  and  oxygen  (^10). 

5.  Solubilities. — a. — Metal. — Gold  is  not  at  all  tarmsnea  or  in  any  way  acted 
upon  by  water  at  any  temperature,  or  by  hydrosulphuric  acid.     Neither  nitric 
nor  hydrochloric  acid  attacks  it  under  any  conditions;  but  it  is  rapidly  attacked 
by  chlorine  (as  gas  or  in  water  solution),  dissolving  promptly  in  nitrohydro- 
chloric acid,  as  auric  chloride,  AuCl3;  by  bromine,  dissolving  in  bromine  water, 
as  aurif  bromide,  AuBr3;  and  by  iodine;  dissolving  when  finely  divided  in  hydri- 
odic  acid  by   aid  of  the  air  and  potassium  iodide,  as  potassium  auric  iodide, 
KIAuI3:  4Au  +  12HI  -f  4KI  +  302  =  4KIAuI3  +  6H20  .     Potassium  cyanide 
solution,  with   aid   of  the  air,   dissolves  precipitated  gold   as  potassium  auro- 
cyowide,  XAu<CN)s:  4Au  +  8KCN  +  O2  +  2H2O  =  4KAu(CN),  +  4KOH  . 

Gold  is  separated,  from  its  alloys  with  silver  and  base  metals,  by  solution  in 
nitric  acid;  the  gold  being  left  as  a  black-brown  powder— together  with 
platinum  and  oxides  of  antimony  and  tin.  When  the  gold-silver  or  gold-copper 
has  not  over  20  per  cent  gold,  nitric  acid  of  20  per  cent  disintegrates  the  alloy, 
and  effects  the  separation;  when  the  gold  is  over  25  per  cent,  silver  or  lead 
[three  parts)  must  be  added,  by  fusion,  to  the  alloy  before  solution.  (If  gold- 
silver  alloy  contains  60  per  cent  or  more  of  silver,  it  is  silver  color;  if  30  per 
cent  silver,  a  light  brass  color;  if  2  per  cent  silver,  it  is  brass  color.) 

If  gold  and  other  metals  are  obtained  in  solution  by  nitrohydrochloric  acid, 
leaving  most  of  the  silver  as  a  residue,  the  noble  metals  can  be  precipitated  by 
zinc  or  ferrous  sulphate,  and  the  precipitate  of  gold,  silver,  etc.,  treated  with 


&2  GOLD.  §73,  56. 

nitric  acid,  which  will  now  dissolve  out  any  proportion  of  silver  not  less  than 
15  per  cent,  to  85  per  cent  of  gold,  and  dissolve  the  baser  metals.  Concentrated 
sulphuric  acid  dissolves  silver,  and  leaves  gold. 

6. — The  oxides  and  hydroxides  of  gold  are  insoluble  in  water,  soluble  in  acids. 
c. — The  salts  of  the  oxyacids  are  not  stable,  being  decomposed  by  hot  water. 
Gold  sulphide  is  insoluble  in  water  or  acids,  except  nitrohydrochloric  acid, 
soluble  in  alkali  sulphides.  Aurous  salts  are  decomposed  by  water,  forming 
Au°  and  Au'" .  Auric  chloride  is  deliquescent;  both  the  chloride  and  bromide 
are  readily  soluble  in  water.  The  iodide  is  decomposed  by  water,  forming 
aurous  iodide.  The  double  chlorides,  bromides,  iodides  and  cyanides  are  soluble 
in  water. 

6.  Reactions,  a.  The  fixed  alkali  hydroxides  and  carbonates  in  excess 
do  not  precipitate  AuCl3  solutions,  as  a  soluble  aurate,  KAu02 ,  readily 
forms;  but  upon  boiling  and  neutralizing  the  excess  of  alkali,  Au(OH).< 
is  precipitated.  Ammonium  hydroxide  precipitates  from  concentrated 
solutions  a  reddish-yellow  ammonium  aurate,  (NH3)2Au203 ,  "  fulminating 
gold."  &.  Oxalic  acid  reduces  gold  chloride  from  solutions,  slowly  (nitric 
acid  should  be  absent  and  the  presence  of  ammonium  oxalate  is  advan- 
tageous), but  completely.  The  gold  separates  in  metallic  flakes  or  forms 
a  mirror  on  the  side  of  the  test-tube.  2AuCl3  +  3H2C204  =  2Au  +  6C02 
+  6HC1 .  As  platinum,  palladium,  and  other  second  group  metals  are 
not  reduced  by  oxalic  acid,  this  method  of  removal  of  gold  should  be 
employed  upon  the  original  solution  before  the  precipitation  of  the  second 
group  metals  as  sulphides.  Potassium  gold  cyanide,  KCN.Au(CN)3 ,  is 
formed  when  a  neutral  solution  of  AuCl3  is  added  to  a  hot  saturated 
solution  of  KCN .  It  is  very  soluble  in  water  and  by  heating  above  200° 
it  is  decomposed  into  CN  and  KCN.AuCN ,  which  latter  product  is  formed 
when  gold  is  dissolved  in  KCN  in  the  presence  of  air  (5a).  c.  A  solution 
of  AuCl3  is  precipitated  as  Au°  by  a  solution  of  KN02 .  d.  Sodium 
pyrophosphate  forms  with  AuCl3  a  double  salt  which  has  found  application 
in  gold  plating,  e.  Hydrosulphuric  acid  precipitates  from  gold  chloride 
solution,  hot  or  cold,  gold  sulphide,  variable  from  Au2S  to  Au2S2 ,  brown, 
insoluble  in  acids,  hot  or  cold,  except  in  nitrohydrochloric  acid,  in  which 
it  readily  dissolves;  soluble  in  alkali  sulphides  to  a  thio-salt.  Alkali 
sulphites  precipitate  gold  chloride  solution  as  double  sulphite,  i.  e. 
Au2(S03)3.(NH4)2S03.6NHs  +  3H20  .  Upon  boiling  the  sulphite  acts  as 
a  reducing  agent,  giving  metallic  gold. 

/.  Potassium  iodide,  added  in  small  portions  to  solution  of  auric  chloride 
(so  that  the  latter  is  constantly  in  excess  where  the  two  salts  are  in 
contact),  and  when  equivalent  proportions  have  been  reached,  gives  a  yel- 
low precipitate  of  aurous  iodide,  Aul ,  insoluble  in  water,  soluble  in  large 
excess  of  the  reagent;  the  precipitate  accompanied  with  separation  of  free 
iodine,  brown,  which  is  quickly  soluble  in  small  excess  of  the  reagent  as  a 
colored  solution:  AuCl3  -f  4KI  =  Aul  +  3KC1  +  I2  with  KI .  But,  on 
gradually  adding  auric  chloride  to  solution  of  potassium  iodide,  so  that  the 


§74,1.  PLATINUM.  93 

latter  is  in  excess  at  the  point  of  chemical  change,  there  is  first  a  dark- 
green  solution  of  potassio-auric  iodide,  KIAuI3  ;  then  a  dark-green  precipi- 
tate of  auric  iodide,  AuI3 ,  very  unstable,  decomposed  in  pure  water,  more 
quickly  by  boiling;  changed  in  the  air  to  the  yellow  aurous  iodide. 

g.  Stannous  chloride  gives  a  purple  precipitate  containing  the  oxides  of 
tin  with  the  gold,  "  purple  of  Cassius  "  insoluble  in  acids. 

h.  Ferrous  sulphate  is  the  most  common  reagent  for  the  detection  of 
gold,  reducing  all  gold  salts  to  the  metallic  state;  AuCl3  +  3FeS04  — 
Au  +  Fe2(S04)3  +  FeCl3 . 

7.  Ignition. — Gold  is  reduced  from  many  of  its  compounds  by  light,  and  from 
all  of  them  by  heat — its  separation  in  the  dry  way  being-  readily  effected  by 
fusion  with  such  reagents  as  will  make  the  material  fusible.  Very  small  pro- 
portions are  collected  in  alloy  with  lead,  by  fusion;  after  which  the  lead  is 
vaporized  in  "  cupellation  "  (§59,  7). 

8.  Detection. — In  the  dry  way  gold  is  detected  by  fusion  of  the  mineral 
matter  with  lead,  to  the  formation  of  a  "  button  "  which  is  then  ignited 
to  drive  off  the  lead,  leaving  the  gold  and  silver  behind  as  the  metals. 
In  the  wet  way  the  material,  if  not  in  solution,  is  digested  with  nitro- 
hydrochloric  acid  which  dissolves  all  the  gold.     The  excess  of  acid  is  re- 
moved by  evaporation  and  the  gold  is  precipitated  by  oxalic  acid  or  ferrous 
sulphate,  and  identified  by  its  color  and  insolubility  in  acids.     If  the 
gold  be   not  removed  from   the   original  solution   it   is   precipitated   in 
Group  II.  by  H2S ,  passes  into  Division  A  (tin  group)  by  (NH4)2S ,  and  may 
be  detected  in  the  flask  of  the  Marsh  apparatus  by  the  usual  methods. 

9.  Estimation. — Gold  is  always  weighed  in  the  metallic  state,  to  which  form 
it  is  reduced:  (1)  By  ignition  alone  if  it  is  a  salt  containing  no  fixed  acid;  if  in 
an  ore,  by  mixing  with  lead  and  fusion  to  an  alloy,  and  final  removal  of  the 
lead  by  ignition  at  a  white  heat  in  presence  of  air.     (2)  By  adding  to  the  solu- 
tion some  reducing  agent,  usually  FeSO4  ,  H2C204  ,  chloral  hydrate,  or  some 
easily  oxidized  metal,  such  as  Zn  ,  Cd  ,  or  Mg  .     (3)  Gold  is  also  estimated  volu- 
metrically  by  H2C2O4  and  the  excess  of  H2C2O4  used,  determined  by  KMnO4  . 

10.  Oxidation. — Gold  is  reduced  to  the  metallic  state  by  very  many 
reducing  agents,  among  which  may  be  mentioned  the  following :  Pb ,  Ag , 
Hg,  Hg',  Sn,  Sn",  As,  As'",  AsH3 ,  Sb ,  Sb'",  SbH3 ,  Bi,  Cu,  Cu', 
Pd,  Pt,  Te,  Fe,  Fe",  Al,  Co,  Ni,  Cr'",  Zn,  Mg,  H2C204 ,  HN02 ,  P, 
H3P02 ,  H3P03 ,  PH3 ,  H2S03 ,  and  a  great  number  of  organic  substances. 


§74.  Platinum.    Pt.  =  195.2  .     Valence  two  and  four. 

1.  Properties. — Specific  gravity  at  17.6°,  21.48  (Deville  and  Debray,  C.  r. 
1860,  60,  1038).  Melting  point,  1755°  (Cr.  B.  S.,  35,  1915).  Pure  platinum  is 
a  tin-white  metal,  softer  than  silver,  hardened  by  the  presence  of  other  metals, 
especially  iridium,  which  it  frequently  contains.  It  is  surpassed  in  ductility 
and  malleability  only  by  Au  and  Ag.  Platinum  black  is  the  finely  divided 
metal,  a  black  powder,  obtained  by  reducing  an  alkaline  solution  of  the  platinous 
salt  with  alcohol  (Low,  B.,  1890,  23,  289);  platinum  sponge,  a  gray  spongy  mass, 
by  ignition  of  the  platinum  ammonium  double  chloride;  platinized  asbestos 
(usually  10  per  cent  Pt),  the  metal  in  finely  divided  form  deposited  by  reduction 


94  PLATINUM.  §74,  2. 

from  the  salt  upon  asbestos.  These  finely  divided  forms  of  platinum  have  great 
power  of  condensation  of  gases,  and  by  their  presence  alone  bring  about  a  num- 
ber of  important  chemical  reactions  (catalytic  reaction);  e.  g.,  a  current  of 
hydrogen  mixed  with  air  ignites  when  passed  over  platinum  black,  also  hydro- 
gen and  chlorine  unite.  SO2  unites  with  O  to  form  SO3  ;  alcohol  is  oxidized  to  acetic 
acid,  formic  and  oxalic  acids  to  CO2  ,  AS'"  to  Asv  ,  etc. 

2.  Occurrence. — Found  in  nature  in  the  metallic  state,  generally  alloyed  with 
palladium,    iridium,    osmium,    rhodium,    ruthenium,    etc.     The    Ural    Mountains 
furnish    the    largest    supply    of    platinum.     The    only    known    native    compound 
is  sperrylite,  PtAs2  ,  found  in  Ontario,  Canada,  and  Macon  Co.,  N.  Car. 

3.  Preparation. — Usually  by  the  wet  method.  The  finely  divided  ore  is  treated 
with  nitrohydrochloric  acid  until  the  platinum  is  all  dissolved.     The  filtrate  is 
then   treated  with   lime   water   to  a  slightly  acid   reaction;    this   removes   the 
greater  part  of  the  Fe  ,  Cu  ,  Ir  ,  Rh  ,  and  a  portion  of  the  Pd  .     The  filtrate  is 
now  evaporated  to  dryness,  ignited  and  washed  with  water  and  hydrochloric 
acid.     This  gives  a  commercial  platinum  which  is  melted   with   six  times  its 
weight  of  lead  and  the  finely  divided  alloy  digested  with  dilute  HNO3  ,  which 
dissolves  out  the  Pb  ,  Cu  ,  Pd  ,  and  Rh  .     The  black  powder  which  remains  is 
dissolved  in  nitrohydrochloric  acid,  the  Pb  remaining,  removed  with  H:S04  , 
and  the  Pt  precipitated  with  NH4C1 .     The  precipitate  contains  a  little  rhodium, 
which  is  removed  by  gently  igniting  the  mass  with  potassium  and  ammonium 
di-sulphate,    and    exhausting    with    wrater,    wrhich    dissolves    out    the    rhodium 
sulphate   (§105,  7).     In  the  laboratory  the  platinum  residues  are  boiled  with 
KOH  or  K,CO3  and  reduced  with  alcohol.     The  fine  black  powder  is  filtered, 
\vashed  with  water  and  hydrochloric  acid  and  ignited. 

4.  Oxides   and   Hydroxides. — Platinum    forms   two   oxides,    PtO    and    Pt02  . 
Platinous  hydroxide  is  formed  by  treating  a  dilute  solution  of  platmous  potas- 
sium chloride  with  NaOH  and  boiling   (Jorgensen,  J.  pr.,   1877,   (2),   16,   344). 
A  black  powder  easily  soluble  in  HC1  or  HBr  ,  reduced  by  formic  acid  to  Pt°  , 
gentle  heating  changes  it  to  the  oxide  PtO.    Platinic  hydroxide,  Pt(OH)4  ,  is 
formed  by  treating  a  solution  of  H2PtCl6  with  Na*CO3  in  excess,  evaporating 
to  dryness,  washing  with  water  and  then  with  acetic  acid.     It  is  a  red-brown 
powder,  soluble  in  NaOH,  HC1 ,  HNO3  ,  and  H2SO4;   insoluble  in   HCSH3O,  . 
Gentle  heating  changes  it  to  the  oxide  PtO2  (Topsoe,  B.,  1870,  3,  462). 

5.  Solubilities. — a— Metal. — Platinum  is  not  affected  by  air  or  water,  at  any 
temperature;  is  not  sensibly  tarnished  by  hydrosulphuric  acid  gas  or  solution; 
and  is  not  attacked  at  any  temperature  by  nitric  acid,  hydrochloric  acid  or 
sulphuric   acid,   but  dissolves  in  nitrohydrochloric   acid  less  readily    than   gold. 
The  substance  obtained  by  evaporating  an  aqua  regia  solution  of  platinum  is 
platinic  chloride  plus  2  molecules  of  hydrochloric  acid,  PtCl4.2HCl  .     This   salt 
is  frequently  called  platinic  chloride  and  its  water  solution  is  the  platinic  chloride 
reagent.     It  is  more  correctly  called  chlorpl  dinic  acid  and  the  formula  written 
H2PtCl6.     b. — Oxides   and   hydroxides. — See   4.     c. — Salts. — Platinum    forms    two 
classes  of  salts  (both  haloid  and  oxy),  platinous  and  platinic.     The  oxysalts  are 
not  stable.     None  of  the  platinous  salts  are  permanently  soluble  in  pure  water. 
The  chloride  is  soluble ^  in  dilute  hydrochloric  acid  and  the  sulphate  in  dilute 
sulphuric  acid.     Platinic  chloride,   PtCl4 ,   and  bromide,   all  the  platinicyanides 
(as  PbPt(CN)e),  and  the  platinocyanides  of  the  metals  of  the  alkalis  and  alkaline 
earths  (as  K2Pt(CN)4),  are  soluble  in  water.     The  platinous  and  platinic  nitrates 
are  soluble  in  water,  but  easily  decomposed  by  it,  with  the  precipitation  of  basic 
salts.     The  larger  number  of  the  metallo-platinic  chlorides  or  "chloroplatinates" 
are  soluble  in  water,  including  those  with  sodium  [Na^PtCle  or  (NaCl)2PtCl4],  barium, 
strontium,  magnesium,  zinc,  aluminum,  copper;    and  those  with  potassium,  and 
ammonium,  are  sparingly  soluble  in  water,  and  owe  their  analytical  importance 
as  complete  precipitates  to    their  insolubility  in  alcohol.     Of  the  metallo-platinous 
chlorides  (the  "chloroplatinites") — those  with  sodium  [Na^PtCLi],  and  barium,  are 
soluble;   zinc,  potassium  and  ammonium,  sparingly  soluble;  lead  and  silver,  insol- 
uble in  water.     Platinic  sulphate,  Pt(SO4)2  ,  is  soluble  in  water  (§10). 

6.  Reactions.  —  a.  —  Platinous    chloride,    PtCl2 ,   is    precipitated  by   KOH    as 
Pt(OH)2 ,    soluble  in  excess  of  the  reagent  to  K2PtO2 ,  potassium  platinite,  which 
solution  is  reduced    by  alcohol    to     "platinum    black       (1).      Platinic    chloride, 
HoPtCle,  a  brown-red  solid,  soluble  in  alcohol  and  water,  forms  with  KOH  or 
NH4OH ,  not  too  dilute,  a  yellow  crystalline  precipitate  of  an  alkali  (K  or  NH4), 


§74,  7,  1.  PLATINUM.  95 

chlorplatinate,  e.  g.,  K2PtCl6,  sparingly  soluble  in  water,  soluble  in  excess 
of  the  alkalis  and  reprecipitated  by  hydrochloric  acid.  K2C03  and  (NH4)2CO3 
give  the  same  precipitate,  insoluble  in  excess  of  the  reagent,  A  more  complete 
precipitation  of  the  K  or  NH,  is  obtained  by  the  use  of  the  chlorides.  The 
sodium  platinum  chloride,  Na.PtCl,,  ,  is  very  soluble  in  water  and  is  not  formed 
by  precipitation  with  sodium  salts.  6. — Oxalic  acid  does  not  reduce  platinum 
salts  (distinction  from  gold).  A  solution  of  chloral  hydrate  precipitates  pla- 
tinum from  its  solutions.  I  Ma  ii  nous  and  platinic  salts  form  with  cyanides  a 
great  number  of  double  salts,  c.— See  5c.  d. — Hypophosphorous  acid  reduces 
platinum  salts  to  metallic  platinum.  Phosphates  do  not  precipitate  platinum 
salts. 

G.  Hydrosulphuric  acid  precipitates  solutions  of  the  platinous  salts  as 
the  black  sulphide,  PtS ,  insoluble  in  acids,  sparingly  soluble  in  water  and 
in  alkali  sulphides;  platinic  salts  are  precipitated  as  platinic  sulphide, 
PtS, ,  black;  slowly  soluble  in  alkali  sulphides  (Kibau,  C.  r.,  1877,  85,  283), 
insoluble  in  acids  except  nitrohydrochloric.  Sulphur  dioxide  decolors  a 
solution  of  platinum  chloride  giving  a  compound  which  does  not  respond 
to  the  usual  reagents  for  platinum  and  requires  long  boiling  with  HC1  for 
the  removal  of  the  S02  (Birnbaum,  A.,  1871,  159,  116). 

/.  The  chlorides  of  potassium  and  ammonium  are  estimated  quantita- 
tively by  precipitation  from  their  concentrated  solutions  with  a' solution 
of  platinic  chloride.  Potassium  iodide  colors  a  solution  of  platinum 
chloride  brown-red  mid  precipitates  the  black  platinic  iodide,  PtI4 ,  excess 
of  the  KI  forming  K.,PtIfl ,  brown,  sparingly  soluble  (5c).  g.  Stannous 
chloride  does  not  precipitate  the  platinum  from  platinic  chloride  (distinc- 
tion from  gold),  but  reduces  it  to  platinous  chloride. 

h.  Ferrous  sulphate  solution  on  boiling  with  a  platinum  chloride  solu- 
tion precipitates  the  platinum  as  the  metal,  the  presence  of  acids  hinders 
the  reduction. 

7.  Ignition. — All  platinum  compounds  upon  ignition  are  reduced  to  the 
metal.  Owing  to  the  high  point  of  fusibility  of  the  metal  and  to  the 
difficulty  with  which  it  is  attacked  by  most  chemicals,  platinum  has 
an  extended  use  in  the  chemical  laboratory  for  evaporating  dishes,  cruci- 
bles, foil,  wire,  etc.  IN  THE  USE  OF  PLATINUM  APPAEATUS  WITHOUT 

IN  NECESSARY  INJURY  IT  SHOULD  BE  REMEMBERED  I 

(1)  That  free  chlorine  and  bromine  attack  platinum  at  ordinary  tem- 
peratures (forming  platinic  chloride,  bromide);  and  free  sulphur,  phos- 
phorus, arsenic,  selenium,  and  iodine,  attack  ignited  platinum  (forming 
platinous  sulphide,  platinic  phosphide,  platinum-arsenic  alloy,  platinic 
selenide,  iodide).  Hence,  the  fusion  of  sulphides,  sulphates,  and  phos- 
phates, with  reducing  agents,  is  detrimental  or  fatal  to  platinum  crucibles. 
The  ignition  of  organic  substances  containing  phosphates  acts  as  free 
phosphorus,  in  a  slight  degree. 

The  heating  of  ferric  chloride,  and  the  fusion  of  bromides,  and  iodides, 
act  to  some  extent  on  platinum. 


96  PLATINUM.  §74,  7, 2. 

(2)  The  alkali  hydroxides  (not  their  carbonates)  and  the  alkaline  earths, 
especially  baryta  and  lithia,  with  ignited  platinum  in  the  air,  gradually 
corrode  platinum   (by  formation  of  platinites:   2Pt   +   2BaO   -}-   02   = 
2BaPtOt .          Silver  crucibles   are  recommended  for  fusion  with   alkali 
hydroxides. 

(3)  All  metals  which  may  he  reduced  in  the  fusion — especially  compounds 
of  lead,  bismuth,  tin,  and  other  metals  easily  reduced  and  melted — and  all 
metallic  compounds  with  reducing  agents  (including  even  alkalis  and  earths) 
form  fusible  alloys  with  ignited  platinum.     Mercury,  lead,  bismuth,  tin, 
antimony,  zinc,  etc.,  are  liable  to  be  rapidly  reduced,  and  immediately  to 
melt  away  platium  in  contact  with  them. 

(-4)  Silica  with  charcoal  (by  formation  of  silicide  of  platinum)  corrodes 
ignited  platinum,  though  very  slowly.  Therefore,  platinum  crucibles 
should  not  be  supported  on  charcoal  in  the  furnace,  but  in  a  bed  of  mag- 
nesia, in  an  outer  crucible  of  clay.  Over  the  flame,  the  best  support  is  the 
triangle  of  platinum  wire. 

(5)  The  tarnish  of  the  gas-flame  increases  far  more  rapidly  upon  the 
already  tarnished  surface  of  platinum — going  on  to  corrosion  and  crack- 
ing. The  surface  should  be  kept  polished — preferably  by  gentle  rubbing 
with  moist  sea-sand  (the  grains  of  which  are  perfectly  rounded,  and  do  not 
scratch  the  metal).  Platinum  surfaces  are  also  cleansed  by  fusing  borax 
upon  them,  and  by  digestion  with  nitric  acid. 

8.  Detection.  —  Platinum  is  identified  by  the  appearance  of  the  reduce** 
metal;  by  its  insolubility  in  HC1  or  HN03  and  solubility  in  HN03  +  HC1  ; 
and  by  its  formation  of  precipitates  with  ammonium  and  potassium 
chlorides  and  KI .  It  is  separated  from  gold  by  boiling  with  oxalic  acid  and 
ammonium  oxalate,  which  precipitate  the  gold,  leaving  the  platinum  in  solu- 
tion. The  filtrate  from  the  gold  should  be  evaporated,  ignited,  and  the 
residue  examined  and  after  proving  insolubility  in  HC1  or  HN03 ,  dissolved 
in  HN08  and  HC1  and  the  presence  of  platinum  confirmed  with 
NH4C1 .  If  the  gold  and  platinum  have  been  precipitated  in  the  second 
group  with  H2S  and  dissolved  with  (NH4)2SX  they  may  be  separated  from 
As ,  Sb  ,  and  Sn  by  dissolving  the  reprecipitated  sulphides  in  HC1  +  KC103 , 
evaporating  to  remove  the  chlorine  and  boiling  after  adding  KOH  in  ex- 
cess, with  chloral  hydrate,  which  precipitates  the  An  and  Pt ,  leaving  the 
As ,  Sb ,  and  Sn  in  solution.  The  An  and  Pt  may  then  be  dissolved  in 
HN03  -|-  HC1  and  separated  as  directed  above.  FeS04  may  be  use  to  pre- 
cipitate An  and  Pt ,  separating  them  from  As ,  Sb ,  and  Sn  . 

9.  Estimation. — Platinum  is  invariably  weighed  in  the  metallic  state.  It  is 
brought  to  this  condition:  (1)  By  simple  ignition;  (2)  by  precipitation  as 
(NH4)2PtCl6  ,  K2PtCl6  ,  or  PtS2  and  ignition;  (3)  by  reduction,  using  Zn  ,  Mg  , 
or  F«SO4  . 

40.  Oxidation. — Solutions  of  platinum  are  reduced  to  the  metallic  state  by  the 


§75,  (k                                                MOLYBDENUM.  9t 

following-   metals:    Pb  ,  Ag  ,    Hg ,    Sn    (Sn"   to    Pt"   only),    Bi ,    Cu  ,  Cd  ,   Zn  , 

Fe  ,    Fe"  ,    Co  ,    and    Ni  .     Very    many    organic    substances    reduce  platinum 
compounds  to  the  metallic  state. 


§75.  Molybdenum.     Mo  =  96.0  .     Valence  two,  three,  four  and  six. 

1.  Properties.—  Specific   gravity,    9.01    (Moissan,    C.    R.,    120,    1895).    Meltifj 
point,   2500°?   (Cr.   B.  S.,  36,    1915).     It  is  a  silver-white,   hard,   brittle  metal, 
not  oxidized  in  the  air  or  water  at  ordinary  temperatures.     Upon  heating  in 
the  air  it  becomes  brown,   then  blue,   and  finally  burns  to  the  white  MoO;  . 
Heated  to   a  red  heat  in  contact  with  steam,  it  forms  first  a  blue  oxide,  then 
MoO3  . 

2.  Occurrence.  —  Not    found    native,     but    occurs   chiefly   as    molybdenite, 
MoS2;  as  an  oxide  in  molybdenum  ochre,  Imolybdite,  MoCX  ;    and  as  wulfenite 
PbMoO4. 

3.  Preparation. — (1)  By  heating  the  oxide,  sulphide  or  chloride  in  a  current 
of   oxygen   free   hydrogen    (von    der  Pfordten,    B.,   1884,    17,    732;    Rogers    and 
Mitchell,  J.  Am.  Soc.,  1900,  22,  350);  (2)  by  heating  with  C  and  Na2CO8;  (3)  by 
heating  Mo03  with  KCN  (Loughlien,  I.e.). 

4.  Oxides  and  Hydroxides. — Molybdous  hydroxide,  MoO.xH^O  ,  is  formed  when 
molybdous  chloride  or  nitrate  is  precipitated  with  alkali  hydroxides  or  carbon- 
ates,  dark  brown  becoming  blue  in  the  air  by  oxidation.     Mo(OH)8  ,   black, 
turning  red-brown  by  oxidation  in  the  air,  is  formed  by  treating  MoCl3  with 
KOH;  also  by  electrolysis  of  ammonium  molybdate  (Smith,  B.,  1880,  13,  751). 
By  heating  the  hydroxide  in   a  vacuum  Mo2Og   is  obtained  as  a  black  mass, 
insoluble  in  acids.     MoOa  ,  a  dark  bluish  mass,  insoluble  in  KOH  or  HC1 ,  is 
formed  by  igniting  a  mixture  of  ammonium  molybdate,  potassium  carbonate 
and  boric  acid,  and  exhausting  the  fused  mass  with  water  (Muthmann,  A.,  1887, 
238,    114).    Molybdic  anhydride    (acid),    MoOs  ,   white,   occurs   in   nature;   it    is 
obtained  by  the  ignition  of  the  lower  oxidized  compounds  in  the  air  or  in  the 
presence  of  oxidizing  agents. 

5.  Solubilities. — Molybdenum  is  readily  soluble  in  nitric  acid  with  oxidation 
to  MoO8  ,  evolving  NO;  in  hot  concentrated  sulphuric  acid,  evolving  SO2  .     The 
various  lower  oxides  of  molybdenum  are  soluble  in  acids  forming  corresponding 
salts,  not  very  stable,  oxidizing  on  exposure,  to  molybdic  acid  and  molybdates; 
on  the  other  hand,  reducing  agents  reduce  molybdates  to  the  lower  forms  of 
molybdenum  salts,  nea»ly  all  of  which  are  colored  brown  to  reddish  brown  or 
violet.     The  salts  of  molybdenum  are  nearly  all  soluble  in  water.     Molybdic 
anhydride,   Mo08  ,   white,   is   sparingly   soluble   in   water  and    possesses   basic 
properties    towards    stronger   acids,    dissolving   in    them    to    form    salts.     The 
chlorides  and  the  sulphates  are  soluble  in  water  (Schulz-Sellack,  B.,  1871,  4,  14); 
the  nitrates  in   dilute   nitric  acid.     The  anhydride   Mo08    combines   with   the 
alkalis  to  form  molybdates,  soluble  in  water.     Molybdates  of  the  other  metals 
are  insoluble  in  water.     Solutions  of  the  alkali  molybdates  are  decomposed  by 
acids  forming,  MoO3  ,  which  dissolves  in  excess  of  the  acids. 

6.  Reactions. — a. — The  dyad,  triad  and  tetrad  molybdenum  salts  are  precipi- 
tated by  the  alkali  hydroxides   and   carbonates,   forming  the   corresponding 
hj^droxides,  insoluble  in  excess  of  the  precipitant.     These  hydroxides  oxidize 
in  the  air  to  a  blue  molybdenum  molybdate.     6. — A  solution  of  a  molybdate 
acidulated  with  hydrochloric  acid  gives  no  red  color  with  KCNS   (distinction 
from  Fe"');  but  if  Zn  be  added,  reduction  to  a  lower  oxide  of  molybdenum 
takes  place  and  an  intense  red  color  is  produced.     Phosphoric  acid  does  not 
destroy   the  color   (difference   from  ferric  thiocyanate).     Upon   shaking  with 
ether  the  sulphocyanate  is  dissolved  in  the  ether,  transferring  the  red  color 
to  the  ether  layer.     In  molybdic  acid  solutions,  acidulated  with  hydrochloric 
acid,  potassium  ferrocyanide  gives  a  reddish  brown  precipitate.     An  alkaline 
solution  of  molybdates  is  colored  a  deep  red  to  brown  by  a  solution  of  tannic 
acid.    <?.— See  5, 


5IUFORNIA  COLLEtf 


98  MOLYBDENUM.  §75,  Qd. 

d  — Tribasic  phosphoric  acid  and  its  salts  precipitate,  from  strong  nitric 
acid  solutions  of  ammonium  molybdate,*  somewhat  slowly,  more  rapidly 
on  warming,  ammonium  pJwspJw-molybdate  ( (NH4)3POi.12Mo03.3H20), 
yellow,  of  variable  composition,  soluble  in  ammonium  hydroxide  and  other 
alkalis,  sparingly  soluble  in  excess  of  the  phosphate.  The  sodium  phospJio- 
molybdate  is  soluble  in  water,  and  precipitates  ammonium  from  its  salts; 
also,  it  "orecipitates  the  alkaloids — for  which  reaction  it  has  some  im- 
portance as  a  reagent.t  Arsenic  acid  and  arsenates  give  the  same  reaction ; 
ammonium  arseno-molybdate  being  formed  (g). 

e. — Neutral  or  alkaline  solutions  of  molybdates  are  colored  yellow  to 
brown  by  hydrosulphuric  acid  but  are  not  precipitated.  From  the  acid 
solutions  a  small  amount  of  the  hydrogen  sulphide  gives  no  precipitate 
but  colors  the  solution  blue ;  with  more  hydrosulphuric  acid  the  brown  or 
red-brown  precipitate,  MoS3 ,  molybdenum  trisulphide,  is  obtained  after 
some  time.  The  precipitate  is  soluble  in  ammonium  sulphide,  better  when 
hot  and  not  too  concentrated,  as  ammonium  thiomolybdate,  (NH4)2MoS4 . 
from  which  acids  precipitate  the  trisulphide  (Berzelius,  Pogg.,  1826,  7, 
429),  soluble  in  nitric  acid,  insoluble  in  boiling  solution  of  oxalic  acid 
(separation  from  stannic  sulphide). 

If  Na2S2O3  be  added  to  a  solution  of  ammonium  molybdate,  slightly  acid, 
a  blue  precipitate  and  blue-colored  solution  is  obtained.  If  the  solution  be 
more  strong^  acid,  a  red  brown  precipitate  is  obtained.  An  acid  solution  of  a 
molybdate  treated  with  hypophosphorous  and  sulphurous  acids  gives  an  in- 
tense bluish  green  precipitate  or  color,  depending  upon  the  amount  of  molyb- 
denum present. 

f. — Halogen  compounds  not  important  in  analysis  of  molybdenum. 

g. — Arsenic  acid,  and  arsenates  form,  with  a  nitric  acid  solution  of  ammonium 
molybdate,  a  yellow  precipitate  of  aininoiiiinn  arseno-motybdate,  in  appearance 
and  reactions  not  to  be  distinguished  from  the  ammonium  phospho-molybdate; 
except  the  precipitation  does  not  take  place  until  the  solutions  are  slightly 
warmed,  Avhile  with  phosphates  the  precipitation  begins  even  in  the  cold. 
Stannous  salts  give  with  (NH4)oMo04  a  blue  solution  of  the  lower  oxides  of 
molybdenum  (a  delicate  test  for  Sn")  (Longstaff,  C.  A'.,  1899,  79,  282). 

hm — The  alkali  molybdates  are  soluble  in  water  and  their  solutions  precipitate 
solutions  of  nearly  all  other  metallic  salts,  forming  molybdates  of  the  corre- 
sponding metals,  insoluble  in  water,  e.g.,  K,MoO*  +  Pb(NO3)2  =  PbMoO4  + 
2KNOS  . 

*  The  reageift  ammonium  molybdate,  (NH4)2  MoO4,  is  prepared  by  dissolving  molybdic  acid, 
MoOs  ( 100  grams],  in  ammonium  hydroxide  (250  cc.  sp.  gr.  0.90  with  250  cc.  water)  cooling,  and 
slowly  pouring  this  solution  into  well  cooled  fairly  concentrated  nitric  acid  (750  cc.  sp.  gr.  1.42 
with  750  cc.  Witter;  with  constant  stirring. 

t  Sodium  Plinsplin-moliilHlatc— Sonnenschein's  reagent  for  acid  solutions  of  alkaloids—  is  pre- 
pared as  follows :  The  yellow  precipitate  formed  on  mixing  acid  solutions  of  ammonium  molyb- 
date and  sodium  phosphate— the  ammonium  phospho-molybdate— is  well  washed,  suspended  in 
water,  and  heated  with  sodium  carbonate  until  completely  dissolved.  The  solution  is  evapor- 
ated to  dryness,  and  the  residue  gently  ignited  till  all  ammonia  is  expelled,  sodium  being  sub- 
stituted for  ammonium.  If  blackening  occurs,  from  reduction  of  molybdenum,  the  residue  is 
moistened  with  nitric  acid,  and  heated  again.  It  is  then  dissolved  with  water  and  nitric  acid 
to  strong  acidulation ;  the  solution  being  made  ten  parts  to  one  part  of  residue.  It  must  be 
kept  from  contact  with  vapor  of  ammonia,  both  during  the  preparation  and  when  preserved 
for  use. 


£75, 10.  MOLYRiHiM  M.  99 

7.  Ignition.— With  microcosmic  salt,  in  the  outer  blow-pipe  flame,  all  com- 
pounds of  molybdenum  give  a  bead  which  is  greenish  while  hot,  and  colorless 
on  cooling;  in  the  inner  flame,  a  clear  green  bead.  With  borax,  in  the  outer 
flame,  a,  head,  yellow  while  hot,  and  colorless  on  cooling;  in  the  inner  flame,  a 
brown  bead,  opaque  if  strongly  saturated  (moly hdous  oxide).  On  charcoal, 
in  the  outer  flame,  molybdic  anhydride  is  vapori/ed  as  a  white  incrustation;  in 
the  inner  flame  (belief  with  sodium  carbonate),  metallic  molybdenum  is 
obtained  as  a  gray  powder,  separated  from  the  mass  by  lixiviation.  Dry  molyb- 
dates,  heated  on  platinum  foil  with  concentrated  sulphuric  acid  to  vaporiza- 
tion of  the  latter  form,  on  cooling  in  the  air,  a  blue  mass. 

8.  Detection. — In  the  ordinary  process  of  analysis,  molybdenum  appears 
in  Division  B  (tin  group)  of  the  second  group  with  As ,  Sb  ,  Sn ,  An ,  and 
Pt  .  The  sulphide  is  insoluble  in  strong  hydrochloric  acid  but  dissolves 
together  with  the  As  and  Pt  in  aqua  regia  or  potassium  chlorate  and  hydro- 
chloric acid.  On  evaporation  the  platinum  is  precipitated  as  potassium 
ehlorplatinate.  From  the  filtrate  the  arsenic  may  be  precipitated  by  means 
of  magnesia  mixture  and  the  gold  by  heating  with  oxalic  acid.  The 
molybdenum  remains  in  solution  as  molybdic  acid.  This  solution, 
evaporated  to  dryness,  dissolved  in  ammonium  hydroxide  and  poured  into 
moderately  concentrated  HC1  forms  a  solution  of  ammonium  molybdate 
which  may  be  identified  by  the  many  precipitation  and  reduction  tests 
(6  &,  c,  d,  e,  i,  etc.,  7,  and  9).  If  the  molybdenum  be  present  as  a  molybdate 
it  may  be  precipitated  from  its  nitric  acid  solution  by  Na^HPO^ ,  washed, 
dissolved  in  ammonium  hydroxide,  the  phosphate  removed  by  magnesia 
mixture  (§189,  6a),  and  the- filtrate  evaporated  to  crystallization  (Maschke, 
Z.,  1873,  12,  380).  The  crystals  may  be  tested  by  the  various  reduction 
tests  for  molybdenum. 

9.  Estimation. — (./)   Molybdic  anhydride   and  ammonium  molybdate  may  be 
reduced  to  the   dioxide  by   heating  in   a  current  of  hydrogen  gas.     The  heat 
must  not  be  permitted  to  rise  above  dull  redness.     Or  the   temperature  man 
rise  to  a  white  heat,  which  reduces  it  to  the  metallic  state,  in  which  form  it  is 
weighed.     (2)   Lead  acetate  is  added  to  the  alkali  molybdate,  the  precipitate 
washed  in  hot  water,  and  after  ignition  weighed  as  PbMoO4  .     (3)   Volumet- 
rically.     The  molybdic  acid  is  treated  with  zinc  and  HC1 ,  which  converts  it  into 
MoCl3  .     This  is  converted  into  molybdic  acid  again  by  standard  solution  of 
potassium  permanganate. 

10.  Oxidation. — Reducing  agents  convert  molybdic  acid  either  into  the  blue 
intermediate  oxides,  or,  by  further  deoxidation,  into  the  black  molybdous  oxide, 
MoO  .     In  the  (hydrochloric)  acid  solutions  of  molybdic  acid,  the  blue  or  black 
oxide  formed  by  reduction,  will  be  held  in  solution  with  a  blue  or  brown  color. 
Nitric    acidulation    is,    of    course,    incompatible    with    the    reduction.     Certain 
reducing  agents  act  as  follows: 

Cane  sugar  in  the.  fVcbly  acid  boiliiig  solution,  forms  the  blue  color — seen 
better  after  dilution;  a  delicate  test,  Stannous  chloride  forms  first  the  blue, 
then  the  brown,  or  the  (j-rccuisli  brown  to  black-brown,  solution  of  both  the 
intermediate  oxide  and  the  molybdous  oxide.  Zinc,  with  HC1  or  H2SO.t  ,  gives 
the  bhie,  then  green,  then  broim  color,  by  progressive  reduction.  Formic  and 
oxalic  acids  do  not  react.  A  solution  of  1  milligram  of  sodium  (or  ammonium) 
molybdate  in  1  cc.  of  concentrated  sulphuric  acid  (about  1  part  to  18-10  parts),  is 
in  use  as  Frcehde's  Reagent  for  alkaloids.  The  molybdenum  in  this  solution, 
which  must  be  freshly  prepared  for  use  each  time,  is  reduced  by  very  many 
organic  substances;  and  with  a  large  number  of  alkaloids,  it  gives  distinctive 
colors,  blue,  red,  brown  and  yellow. 


100  BISMUTH.  §76,  1. 


THE  COPPER  GROUP  (SECOND  GROUP,  DIVISION  B). 

Mercury  (Mercuricum),  Lead,  Bismuth,  Copper,  Cadmium  (Ruthenium, 
Rhodium,  Palladium,  Osmium). 

§76.  Bismuth,  Bi  =  208.0.       Valence  three  and  five. 

1.  Properties.— Specific  gravity,  9.7474  (Classen,  B.,  1890,  23,  938);  melting  point, 
271°  (Cir.  B.  S.,  36,  1915);    it  vaporizes  atfl700°  and  the  density  of  the  vapor 
shows  that  the  molecule  Bi  has  begun  to  dissociate   (Biltz  and  V.    Meyer,   B., 
1889,   22,   725).     It  is  a   hard,   brittle,   reddish-white,   lustrous   metal;    forming 
beautiful  rhombohedral   crystals   when   a   partially   cooled   mass  is   broken   into 
and  the  still  molten  mass  decanted.     Alloys  of  bismuth  with  other  metals  give 
compounds  of  remarkably  low  melting  points,  e.  g.,  an  alloy  of:    Pb  three,  Sn 
four,  Bi  fifteen,  and  Cd  three  by  weight  melts  at  55.5°;    and  an  alloy  of:    Bi 
four.  Pb  two,   Sn  one.   and  Cd  one  parts  by  weight  melts  at  65.5°   "Wood's 
Metal."  * 

2.  Occurrence. — It  is  a  comparatively  rare  metal,  not  very  widely  distributed; 
frequently  found  native.     It  is  found  in  greatest  quantities  in  Saxony:   also  found 
in    Bohemia,  France,  England  and  South  America.     As  mineralogical  varieties  it 
occurs  as  Bismite  (Bi2O3),  bismutite  (4Bi2Oj.3CO2.4H2O),  bismuthinite  (BijSj),  etc. 

3.  Preparation. — The  rock  containing-  bismuth,  usually  with  large  amounts 
of   cobalt,   etc.,   is    roasted   to   remove    sulphur   and    arsenic,    which    is    nearly 
always  present.     The  mass  is  then  fused  with  charcoal.     The  molten  bismuth 
settles  to  the  bottom  below  the  layer  of  cobalt.     The  cobalt  becomes   solid 
while  the  bismuth  is  still  molten,   and  the   two   are   separated   mechanically. 
The  metal  is  further  purified  by  melting  with  KNO8  or  KCN  . 

4.  Oxides. — Bismuth  trioxide,  Bi2O3  ,  is  formed  by  heating  the  metal  in  the 
presence  of  air,  or  by  igniting  the  hydroxide;  it  is  a  pale  citron-yellow  powder. 
The  hydroxide,  Bi(OH),  ,  white,  is  formed  by  precipitating  a  solution  of  a  salt 
of  bismuth  with  an  alkali  hydroxide.     If  bismuth  chloride  is  used  the  hydroxide 
formed  always  contains  some  oxychloride,  BiOCl  (Strohmeyer,  Pogg.,  1832,  26, 
549).  The  meta  hydroxide,  BiO(OH)  ,  is  formed  upon  drying  the  orthohydroxide 
at  100°   (Arppe,  Pogg.,  1845,  64,  237).    Bismuth  pentoxide,  Bi2O5  ,  is  formed  by 
igniting  Bi(OH),  with  excess  of  KOH  or  NaOH  in  presence  of  the  air,  and 
washing  the  cooled  mass  repeatedly  with  cold  dilute  nitric  acid  (Strohmeyer, 
I,  c.);  or  by  treating  Bi(OH)s  with  three  per  cent  H202  in  strong  alkaline  solu- 
tion (Hasebrock,  B.,  1887,  20,  213).     It  is  a  heavy  dark  brown  powder.     At  150° 
it  gives  off  O,  and  at  the  temperature  of  boiling  mercury  becomes  Bi20s  .     It 
is  decomposed  in  the  cold  by  HC1  with  evolution  of  chlorine.     Bismuthic  acid, 
HBiO8  ,   or   more   probably   Bi2O5.H2O ,    is   formed    upon    conducting    a   rapid 
current  of  chlorine   into  Bi(OH)s    suspended  in   concentrated  KOH   solution. 
It  is  a  beautiful  scarlet  red  powder  which  at  120°  gives  off  its  water,  becoming 
Bi205  (Muir,  J.  C.,  1876,  29,  144;  Muir  and  Carnegie,  J.  <7.,  1887,  51,  86).     It  is 
doubtful  if  any  alkali  salt  of  bismuthic  acid  exists,  although  mixtures  of  KBiO3 
and  HBi03  are  claimed  by  Hoffmann  (A.,  1884,  223,  110),  and  Andre  (C.  r.,  1891, 
113,  860).     The  so-called  bismuth  tetroxide,  Bi204  ,  is  probably  a  mixture  of  the 
trioxide  and  pentoxide  (§12). 

5.  Solubilities. — a. — Metal— Metallic    bismuth    is    insoluble    in    hydrochloric 
acid  f;  soluble  in  warm  concentrated  sulphuric  acid  with  evolution  of  sulphur 
dioxide;  readily  soluble  in  nitric  acid  and  in  nitrohydrochloric  acid.     It  burns 
in  chlorine  with  production  of  light;  it  combines  with  bromine,  but  more  slowly 
than  antimony;  it  combines  readily  upon  fusing  together  with  I ,  S  ,  Se  ,  Te  , 
As  ,  and  Sb  ,  besides  the  many  metals  with  which  it  combines  to  form  com- 

*  For  other  fusible  alloys  see  Van  Nostrand's  Chemical  Annual,  3d  issue,  p.  376. 

t  A  trace  of  bismuth  can  always  be  found  in  solution  when  the  metal  is  boiled  with  hydro- 
chloric acid,  but  no  more  than  when  the  metal  has  been  boiled  with  pure  water  (Ditte  and 
MeUner,  A.  Ch.,  1896,  (6),  29,  389), 


§76, 6a.  BISMUTH.  101 

mercial  alloys  (1).  The  halogen  derivatives  of  pentad  bismuth  are  not  known 
(Muir,  J.  C.,  1876,  29,  144).  b — Oxides  and  hydroxides. — Bismuth  oxide,  BiaO3  , 
and  the  hydroxides,  Bi(OH)3  and  BiO(OH),  are  soluble  in  hydrochloric,  nitric 
and  sulphuric  acids;  insoluble  in  water  and  the  alkali  hydroxides  or  carbonates. 
The  presence  of  glycerol  prevents  the  precipitation  of  bismuth  hydroxides 
from  solutions  of  its  salts  by  the  alkalis.*  Bismuth  pentoxide,  Bi205  ,  is  solu- 
ble in  HC1  ,  HBr  ,  and  HI  with  evolution  of  the  corresponding  halogen  and 
formation  of  the  triad  salt.  Nitric  and  sulphuric  acids  in  the  cold  have  but 
little  or  no  action;  when  hot  the  triad  bismuth  salt  is  formed  with  evolution 
of  oxygen. 

c. — Salts. — Most  of  the  salts  of  bismuth  are  insoluble  in  water.  The 
chloride,  bromide,  iodide,  nitrate,  and  sulphate  are  soluble  in  water  acidu- 
lated with  their  respective  acid,  or  with  other  acids  forming  "  soluble  " 
bismuth  salts.  Pure  water  decomposes  the  most  of  the  solutions  of  bis- 
muth salts  forming  corresponding  oxy-salts  (§70,  5d  footnote). 

The  chloride,  bromide  and  sulphate  are  deliquescent. 

d. — Water. — A  solution  of  bismuth  chloride  in  water  acidulated  with 
hydrochloric  acid  is  precipitated  on  further  dilution  with  water,  bismuth 
oxy-chloride,  BiOCl  being  formed;  e.  g.,  BiCl3  +  H20  =  BiOCl  +  2HC1 , 
insoluble  in  tartaric  acid  (distinction  from  antimony,  §70,  5d).  The  hydro- 
chloric acid  set  free  serves  to  hold  a  portion  of  the  bismuth  in  solution. 
The  presence  of  acetic,  citric,  and  other  organic  acids  prevents  the  pre- 
cipitation of  solutions  of  bismuth  salts  upon  further  dilution  with  water. 
The  washing  of  the  precipitated  oxy-salt  with  pure  water  removes  more  of 
the  acid  forming  a  salt  still  more  basic. 

Bi(N03)3  +  HaO  =BiONO,  +  2HNO3 
12BiON08  +  H3O  =  6Bi2O8,5N2O6  +  2HN03 

This  is  prevented  by  the  presence  of  one  part  ammonium  nitrate  to  five 
hundred  parts  water  (Lowe,  /.  pr.,  1858,  74,  341). 

Bismuth  nitrate  crystallizes  with  ten  molecules  of  water,  Bi(N03)3. 
10H20  .  It  is  decomposed  by  a  small  amount  of  water  forming  the  basic 
nitrate,  BiON03  ;  this  is  soluble  in  dilute  nitric  acid,  when  further  dilution 
with  water  to  any  extent  is  possible  without  precipitation  of  the  basic 
salt,  but*  a  drop  of  hydrochloric  acid  or  a  chloride  causes  a  precipitate  of 
the  oxychloride  in  the  diluted  solution.  The  bromide  is  readily  decom- 
posed by  water  to  BiOBr  ;  the  iodide  is  stable  to  cold  water,  but  is  decom- 
posed by  hot  water  to  BiOI  (Schneider,  A.  Ch.,  1857  (3),  50,  488);  the 
normal  sulphate  very  readily  absorbs  water  to  form  Bi2(S04)3.3H20 ,  which 
is  decomposed  by  more  water  to  Bi2O..S03 . 

6.  Reactions,  a. — The  alkali  hydroxides  precipitate  from  solutions  of 
bismuth  salts  lismutli  hydroxide,  Bi(OH)3 ,  white;  insoluble  in  excess  of 
the  fixed  alkalis  (distinction  from  Sb  and  Sn),  insoluble  in  ammonium 

*  Lowe  (C.  2V.,  1882,  45,  296)  dissolves  the  hydroxides  of  copper  and  bismuth  in  glycerol,  adds 
glucose  and  gently  warms.  The  copper  is  completely  precipitated  and  separated  from  the  bis- 
muth. Upon  boiling  the  filtrate  for  some  time  the  bismuth  is  completely  precipitated  as  the 
metal. 


102  BISMUTH.  §76,  65. 

hydroxide  (distinction  from  Cu  and  Cd).  The  hydroxide  is  converted  "by 
boiling  into  the  oxide,  Bi203 ,  yellowish  white.  The  precipitation  is  pre- 
vented by  the  presence  of  tartaric  acid,  citric  acidj  glycerol,  and  certain 
other  organic  substances  (Kohler,  J.  (7.,  1886,  50,  428). 

The  alkali  carbonates  precipitate  basic  bismuth  carbonate,  Bi203.C02  ,  white, 
insoluble  in  excess  of  the  reagent.  Freshly  precipitated  barium  carbonate 
forms  the  same  precipitate  without  heating-. 

&. — Oxalic  acid  and  soluble  oxalates  precipitate  liixnmtJi  o.r<i-l<i1c,  Bi2(C2O4)s  , 
white,  soluble  in  moderately  dilute  acids.  Potassium  cyanide  forms  a  white 
crystalline  precipitate  insoluble  in  excess  of  the  reagent  but  soluble  in  nitric 
or  hydrochloric  acid.  Potassium  ferrocyanide  forms  a  yellowish  white  pre- 
cipitate, potassium  ferricyanide  a  brownish  yellow,  both  soluble  in  hydrochloric 
acid. 

c. — The  action  of  nitric  acid  upon  bismuth  and  its  salts  is  fully  explained 
under  (5).  d. — Metallic  bismuth  is  precipitated  when  bismuth  salts  are  warmed 
with  hypophosphorous  acid  (separation  from  Zn  and  Cd)  (Muthmann  and 
Mawron,  Z.,  1874,  13,  209).  From  solutions  of  bismuth  nitrate  (5d)  phosphoric 
acid  and  soluble  phosphates  precipitate  bismuth  phosphate,  BiPO4  ,  white, 
readily  soluble  in  HC1;  from  solutions  of  the  chloride,  diluted  as  much  as  pos- 
sible without  precipitation,  phosphoric  acid  gives  no  precipitate,  but  the  pre- 
cipitate of  the  phosphate  (soluble  in  HC1)  is  obtained  with  soluble  phosphates. 

e. — Hydrosulphuric  acid  and  sulphides  precipitate  bismuth  sulphide, 
Bi2S3 ,  black,  insoluble  in  dilute  acids  and  in  alkali  hydroxides;  insoluble  in 
alkali  sulphides  (distinction  from  the  metals  of  the  tin  group)  and  in  alkali 
cyanides  (distinction  from  copper).  It  is  soluble  by  moderately  concen- 
trated nitric  acid  (distinction  from  mercury),  the  sulphur  mostly  remain- 
ing free. 

Sodium  thiosulphate  when  warmed  with  solutions  of  bismuth  salts  precipitates 
bixinnlh  sulphide.  Sulphuric  acid  does  not  precipitate  solutions  of  bismuth 
chloride  or  nitrate.  Potassium  sulphate  gives  a  precipitate  with  solutions  of 
both,  that  with  the  chloride  being  apparently  caused  by  the  dilution  of  the 
solution. 

/. — Hydrochloric  acid  and  soluble  chlorides  form  a  precipitate  of  bis- 
muth oxy-chloride,  BiOCl ,  in  solutions  of  bismuth  nitrate  not  containing 
too  much  free  nitric  acid.  This  makes  it  possible  for  bismuth  to  be  precipi- 
tated with  the  silver  group  salts  (§63,  66).  The  precipitate  \s>  readily 
dissolved  on  addition  of  more  hydrochloric  or  nitric  acid  (distinction  from 
the  silver  group  chlorides). 

Hydrobromic  acid  and  soluble  bromides  do  not  precipitate  solutions  of  bis- 
muth chloride,  but  do  precipitate  solutions  of  the  nitrate,  forming  the  oxy- 
bromide,  BiOBr  ,  white.  The  presence  of  potassium  bromide  prevents  the  pre- 
cipitation of  a  bismuth  chloride  solution  by  water  and  also  dissolves  the  oxy- 
chloride  which  has  been  precipitated  by  the  addition  of  water. 

Hydriodic  acid  and  soluble  iodides  precipitate  from  solutions  of  bismuth 
salts,  unless  strongly  acid,  bismuth  iodide,  black  or  brownish  gray  crystals, 
quite  readily  soluble  in  excess  of  the  reagent  *  or  in  strong  HC1  without  warm- 

*  Bismuth  iodide  dissolves  in  solution  of  potassium  iodide  with  an  intense  yellow  color,  deli- 
cate to  one-millionth  (Stone  J.  Sac.  Chem.  Jwl.,  1887,  6,  416).  The  potassium  iodide  solution  of 
bismuth  iodide  is  used  as  Drag-endorff's  reagent  to  detect  the  presence  of  an  alkaloid.  Leger 
(Bf,,  1888,  50,  91)  uses  cmchonine  and  potassium  iodide  to  prove  the  presence  of  bismuth.  Del- 
icate to  one-five  hundred  thousandth.  Other  metals  must  be  removed. 


§76,  9.  BISMUTH.  103 

ing1.  It  is  reprecipitated  on  diluting  the  solution  with  water.  Bismuth  iodide 
is  scarcely  at  all  decomposed  l>.v  washing  with  cold  water,  but  on  boiling  with 
water  it  is  deeom  posed  into  bismuth  o\y-iodide,  BiOI  ,  red,  insoluble  in  KI  , 
soluble  in  HC1  ,  and  in  HI  (Gott  and  Muir,  J.  C'.,  1888,  53,  137). 

Chloric  acid  dissolves  bismuth  hydroxide,  but  the  compound  deeomposes  upon 
evaporation  (\Vachter,  A.,  1844,  52,  233).  Potassium  bromate  and  iodate  both 
precipitate  solutions  of  bismuth  nitrate.  The  iodate  formed  is  scarcely  soluble, 
the  bromate  easily  soluble  in  HN03  . 

g. — Potassium  or  sodium  stannite  hot,  when  added  in  excess  to  bismuth 
solutions,  cause  a  black  precipitate,  from  reduction  to  metallic  bismuth,  a 
very  delicate  reaction.*  The  stannite  is  made,  when  wanted,  by  adding 
to  a  stun  nous  chloride  solution,  in  a  test-tube,  enough  sodium  or  potas- 
sium hydroxide  to  redissolve  the  precipitate  at  first  formed:  2BiCl3  -f- 
3K2Snol  +  (iKOH  =  2Bi  +  6KC1  +  3K,Sn03  +  3H20  (Vanino  and  Treu- 
l.H>rt,  B.,  1898,  31,  1113). 

h. — Solutions  of  bismuth  salts,  nearly  neutral,  poured  into  a  hot  solution  of 
potassium  bichromate  precipitates  the  orange  red  chromate,  (BiO)2Cr2O7;  but 
if  poured  into  a  cold  solution  of  the  neutral  chromate  a  citron-yellow  precipi- 
tate, .iBLCK.L'OO,  ,  is  formed.  These  precipitates  are  soluble  in  moderately 
concentrated  acids,  insoluble  in  fixed  alkalis  (distinction  from  Pb).  The  pre- 
cipitate with  KvCro07  is  used  in  the  quantitative  determination  of  bismuth  (9). 

7.  Ignition. — On  charcoal,  with  sodium  carbonate,  before  the  blow-pipe,  bis- 
muth is  readily  reduced  from  all  its  compounds.     The  (/lobule  is  easily  fusible, 
brittle  (distinction  from  lead),  and  gradually  oxidizable  under  the  flame,  form- 
ing an   Incrustation    (Bi2O3),  orange-yellow  while  ho V "lemon-yellow  when  cold, 
the  edges  bluish-white  when  cold.     The  incrustation  disappears,  or  is  driven 
by  the  reducing  flame,  without  giving  color  to  the  outer  flame.     With  borax 
or  microcosmic  salt,  bismuth  gives  beads,  faintly  yellowish  when  hot,  colorless 
when  cold. 

A  mixture  of  equal  parts  cuprous  iodide  and  sulphur  forms  an  excellent 
reagent  for  the  detection  of  bismuth  in  minerals  by  the  use  of  the  blow-pipe. 
The  reagent  mixed  with  the  unknown  is  fused  on  charcoal  or  on  a  piece  of 
aluminum  sheet.  A  red  sublimate  indicates  bismuth.  Mercury  gives  a  mix- 
ture of  red  and  yellow  sublimates  (Hutchings,  C.  N.,  1877,  36,  249). 

Bismuth  chloride  may  be  sublimed  at  the  temperature  of  boiling  sulphur; 
recommended  as  a  separation  from  lead  (Remmler,  B.,  1891,  24,  3554). 

8.  Detection. — Bismuth  is  precipitated  from  its  solutions  by  H2S  form- 
ing Bi2S3 .     By  its  insolubility  in  (NH4)2S  or  (NHJJS.,  and  solubility  in 
hot  dilute  HN03  it  is  separated  with  Pb ,  Cu  ,  and  Cd  from  the  other  metals 
of  the  tin  and  copper  group.     Dilute  H2S04  removes  the  lead  and  NH,OH 
precipitates  the  bismuth  as  Bi(OH)3  ,  leaving  the  Cu  and  Cd  in  solution. 
The  presence  of  the  bismuth  is  confirmed  by  the  action  of  a  hot  solution  of 
K2Sn02  or  NaOH  and  formaldehyde  f  on  the  white  precipitate  of  Bi(OH)3, 
giving  metallic  bismuth  (6*7)  or  by  dissolving  the  Bi(OH)3  in  HC1  and  its 
precipitation  as  BiOCl  upon  dilution  with  water  (5d). 

9.  Estimation. — (1)    As    metallic    bismuth   formed    by   fusion  with    potassium 
cyanide.      (2)   As  Bi2O3  formed  by  ignition  of  bismuth  salts  of  organic  acids,  or 
of  the  salts  of  volatile  inorganic  oxyacids.      (3}   By  precipitation  by  HsS  ,   and 

*For  a  modification  of  this  test  se3  Muir  (J.  C .,  1S77,  :\",  43). 

t  Sodium  stannite  reduces  lead  hydroxide  while  formaldehyde  does  not.  Traces  of  lead  may 
be  present  owing  to  imperfect  separation. 


104  COPPER.  §76,10. 

after  drying  at  100°,  weighing-  as  Bi2Ss  .  (4)  By  precipitation  by  K2Cr207  ,  and 
after  drying  at  120°,  weighing  as  (BiO)2Cr2O7  .  (.5)  VolumetricaUy.  By"  precipi- 
tation with  K2Cr207  .  Dissolve  the  chromate  in  dilute  acid,  transfer  to  an 
azotometer  and  reduce  the  chromate  with  hydrogen  peroxide  (Baumann,  Z. 
angew.,  1891,  331).  (6)  By  precipitation  as  a  phosphate  with  standard  sodium 
phosphate;  dilution  to  definite  volume  and  determination  of  the  excess  of 
phosphate  in  an  aliquot  part  with  uranium  acetate  (Muir,  J.  (7.,  1877,  32,  674). 

10.  Oxidation. — Metallic  bismuth  reduces  salts  of  Hg?  Ag,  Pt ,  and 
An  to  the  metallic  state.  Bismuth  is  precipitated  as  free  metal  from  its 
solutions  by  Pb ,  Sn ,  Cu ,  Cd ,  Fe  ,  Al ,  Zn  ,  Mg ,  and  HH2P02  (6d).  All 
salts  of  bismuth  are  oxidized  to  Bi205  by  Cl  or  H202  in  strong  alkaline 
mixture  (Hasebrock,  B.,  1887,  20,  213;  Scruff,  A.  Ch.,  1861  (3),  63,  474). 
All  compounds  of  bismuth  are  reduced  to  the  metal  by  potassium  stannite 
K2Sn02  (60).  Bismuth  chloride  or  bromide  heated  in  a  current  of  hydro- 
gen is  partially  reduced  to  the  free  metal  (Muir,  J.  C.,  1876,  29,  144). 
It  is  precipitated  as  free  metal  upon  warming  in  alkaline  mixture  with 
grape  sugar  (56). 


§77.  Copper  (Cuprum)  Cu  =  63.57  .  Valence  one  and  two. 

1.  Properties. — Specific  gravity,  electrolytic,  8.914;  melted,  8.921;  natural  crys- 
tals, 8.94;  rolled  and  hammered  sheet,  8.952  to  8.958   (Marchand  and  Scheerer, 
J.  pr.,  1866,  97,  193).     Melting  point,  1083°  (Cir.  B.  S.,  35,  1915).     A  red  metal, 
but   thin    sheets   transmit   a   greenish-blue   light,    and   it   also   shows   the   same 
greenish-blue  tint  when  in  a  molten  condition.     Of  the  metals  in  ordinary  use, 
only  gold  and  silver  exceed  it  in  malleability.     In  ductility  it  is  inferior  to  iron 
and  cannot  be  so  readily  drawn  into  exceedingly  fine  wire.     Although  it  ranks 
next  to  iron  in  tenacity,  its  wire  bears  about  half  the  weight  which  an  iron  wire 
of  the  same  size  would  support.     As  a  conductor  of  heat  it  is  surpassed  only  by 
gold.     Next  to  silver  it  is  the  best  conductor  of  electricity.     Dry  air  has  no 
action  upon  it;    in  moist  air  it  becomes  coated  with  a  film  of  oxide  which  pro- 
tec  ls  it  from  further  action  of  air  or  of  water.     It  forms  a  number  of  very  im- 
portant alloys  with  other  metals;    bronze   (copper  and  tin),   brass   (copper  and 
zinc  with  sometimes  small  amounts  of  lead  or  tin),  German  silver  (copper,  nickel 
and  zinc). 

2.  Occurrence. — Copper  is  found  native  in  various  parts  of  the  world,  and 
especially  in  the  region  of  Lake  Superior.     It  is  found  chiefly  as  sulphides  in 
enormous  quantities  in  Montana,  Colorado,  Chili  and  Spain;    as  a  carbonate  in 
Arizona.     It    is    very    widely    distributed    and    occurs    in    various    other   forms. 
Chalcopyrite  (CuFeS2);    chalcocite  (Cu2S);   green  malachite  (Cu2(OH)2CO,);   blue 
malachite  (Cu3(OH)2(CO)j);  cuprite  (Cu2O);  and  tenorite  (melanconite)  (CuO). 

3.  Preparation. — For  the  details  of  the  various  methods  of  copper-owning- 
and  refining,  the  works  on  metallurgy  should  be  consulted.     In  the  laboratory 
pure  copper  may  be  produced  (1)  by  electrolysis;  (2)  reduction  by  ignition  in 
hydrogen   gas;    (5)    reduction   of   the   oxide   by   ignition  with   carbon,    carbon 
monoxide,  illuminating  gas,  or  other  forms, of  carbon;    (4)    reduction   of  the 
oxide  Toy  K  or  Na  at  a  temperature  a  little  above  the  melting  point  of  these 
metals;  (5)  reduction  by  fusion  with  potassium  cyanide:  CuO  +  KCN  =  Cu  -f 
KCNO  .     For  its  reduction  in  the  wet  way,  see  10. 

4.  Oxides  and  Hydroxides.— Cuprous  oxide  (Cu2O),  red,  is  found  native;  it  is 
prepared:  (1)  by  reducing  CuO  by  means  of  grape-sugar  in  alkaline  mixture; 
(2)  by  igniting  CuO  with  metallic  copper;  (3)  by  treating  an  ammoniacal  cupric 
solution  with  metallic  copper;  then  adding  KOH  and  drying.     Cuprous  hydrox- 
ide, CuOH  ,  brownish  yellow,  is  formed  by  precipitating  cuprous   salts  with 
KOH  or  NaOH  .     Cupric  oxide,  CuO  ,  black,  is  formed  by  igniting  the  hydroxide. 


§77, 5c.  COPPER.  105 

carbonate,  sulphate,  nitrate  and  some  other  cupric  salts  in  the  air;  or  by 
heating-  the  metal  in  a  current  of  air.  Cupric  hydroxide,  Cu(OH)2  ,  is  formed 
by  precipitating-  cupric  salts  with  KOBE  or  NaOH  .  It  is  stated  by  Rose  (Pogg., 
1863,  120,  1)  that  tetracupric  monoxide,  (Cu4O  ,  is  formed  by  treating1  a  cupric 
salt  with  KOH  and  a  quantity  of  K,SnO,  insufficient  to  reduce  it  to  the  metallic 
state.  A  peroxide  of  copper,  CuO,  ,  is  supposed  to  be  formed  by  treating 
Cu(OH)2  with  H202  at  0°  (Kriiss,  #.,  1884,  17,  2593).  (§10.) 

5.  Solubilities. — a. — Metal. — Copper  does  not  readily  dissolve  in  acids  with 
evolution  of  hydrogen;  it  dissolves  most  readily  in  nitric  acid  chiefly  with 
evolution  of  nitric  oxide'  3Cu  +  8HN03  =  3Cu(N03)2  +  iH2O  -f-  2NO  (Freer 
and  ILigley,  Am..  1899,  21,  377);  also  in  hot  concentrated  sulphuric  acid,  with 
evolution  of  sulphurous  anhydride:  Cu  +  2H2S04  =  CuS04  +  2H2O  -f  SO2  .  If 
dry  hydrochloric  acid  gas  be  passed  over  heated  copper,  CuCl  is  formed  with 
evolution  of  hydrogen  (Weltzien,  A.  Ch.,  18G5,  (4),  6,  487).  A  saturated  solution 
of  hydrochloric  acid  at  15°  dissolves  copper  as  CuCl  with  evolution  of  hydrogen. 
The  action  is  very  rapid  if  the  copper  be  first  immersed  in  a  platinum  chloride 
solution.  Heat  favors  the  reaction  and  the  presence  of  10H.O  to  one  HC1  pre- 
vents the  action  (Engel,  C.  r.,  1895,  121,  528).  Hydrobromic  acid  concentrated 
acts  slowly  in  the  cold  and  rapidly  wrhen  warmed,  forming  CuBr2  ,  with  evolu- 
tion of  hydrogen.  Cold  hydriodic  acid,  in  absence  of  iodine,  is  without  action 
(Mensel,  B.,  1870,  3,  123).  Ammonium  sulphide,  (NH4)2S  ,  colorless,  acts  upon 
copper  turnings  with  evolution  of  hydrogen,  forming  Cu,S  (Heumann,  J.  C.y 

1873,  26,  1105). 

&. — Oxides. — Cuprous  oxide  and  hydroxide  are  insoluble  in  water,  soluble 
in  hydrochloric  acid  with  formation  of  cuprous  chloride,  white,  unstable, 
readily  oxidized  by  the  air  to  colored  cupric  salts.  Cupric  oxide,  black, 
and  hydroxide,  blue,  are  insoluble  in  water,  soluble  in  dilute  acids;  in  a 
mixture  of  equal  parts  glycerine  and  sodium  hydroxide,  sp.  gr.  1.20  (sepa- 
ration from  Cd)  (Donath,  J.  C.,  1879,  36,  178),  in  a  mixture  of  tartrates 
and  fixed  alkalis  (but  precipitated  as  Cu20  by  heating  with  glucose)  (sepa- 
ration from  Cd  and  Zn)  (Warren,  C.  N.,  1891,  63,  193);  insoluble  in 
ammonium  hydroxide  in  absence  of  ammonium  salts  (Maumene,  J.  C., 
1882,  42,  1266). 

c. — Salts. — All  salts  of  copper,  except  the  sulphides,  are  soluble  in  am- 
monium hydroxide.  All  cuprous  salts  are  insoluble  in  water,  soluble  in 
hydrochloric  acid  and  reprecipitated  upon  addition  of  water.  They  are 
readily  oxidized  to  cupric  salts  on  exposure  to  moist  air.  Cuprous  chloride 
and  bromide  are  soluble  in  ammonium  chloride  solution  (Mohr,  J.  C., 

1874,  27,   1099).     Cupric  salts,  in  crystals  or  solution,  have  a  green  or 
blue  color;  the  chloride  (2  aq.)  in  solution  is  emerald-green  when  concen- 
trated, light  blue  when  dilute;  the  sulphate   (5  aq.)  is  "blue  vitriol." 
Anhydrous  cupric  salts  are  white.     The  crystallized  chloride  and  chlorate 
are    deliquescent;    the    sulphate,    permanent;    the    acetate,    efflorescent. 
Cupric   basic   carbonate,   oxalate,   phosphate,   borate,   arsenite,    sulphide, 
cyanide,  ferrocyanide,  ferricyanide,  and  tartrate  are  insoluble  in  water. 
The  ammonio  salts,  the  potassium  and  sodium  cyanides,  and  the  potassium 
and  sodium  tartrate,  are  soluble  in  water.     In  alcohol  the  sulphate  and 
acetate  are  insoluble;  the  chloride  and  nitrate,  soluble.     Ether  dissolves 
the  chloride. 


106  COPPER.  §77,  6a. 

6.  Reactions. — n. — Fixed  alkali  hydroxides  precipitate  acid  solutions  o€ 
cuprous  chloride,  first  as  the  white  cuprous  chloride,  changing  with  more  of 
the  alkali  to  the  yellow  cuprous  hydroxide,  insoluble  in  excess.  Ammonium 
hydroxide  and  carbonate  precipitate  and  redissolve  the  hydroxide  to  a  color- 
less solution,  which  turns  blue  on  exposure.  The  colorless  ammoniacal  solution 
is  precipitated  by  potassium  hydroxide.  Fixed  alkali  carbonates  precipitate 
the  yellow  cuprous  carbonate,  Cu2C03  . 

Fixed  alkalis — KOH — added  to  saturation  in  solutions  of  cupric  salts, 
precipitate  cupric  hydroxide,  Cu(OH)2 ,  deep  blue,  insoluble  in  excess  unless 
concentrated  (Loe\v,  Z.,  1870,  9,  463),  soluble  in  ammonium  hydroxide  (if 
too  much  fixed  alkali  is  not  present),  very  soluble  in  acids,  and  changed, 
by  standing,  to  the  black  compound,  Cu302(OH)2;  by  boiling,  to  CuO  . 
If  tartaric  acid,  citric  acid,  grape-sugar,  milk-sugar,  or  certain  other 
organic  substances  are  present,  the  precipitate  either  does  nqt  form  at  all, 
or  redissolves  in  excess  of  the  fixed  alkali  to  a  blue  solution.  The  alkaline 
tartrate  solution  may  be  boiled  without  change;  in  presence  of  glurnso, 
the  application  of  heat  causes  the  precipitation  of  the  yellow  cuprous 
oxide.  Alkali  hydroxides,  short  of  saturation,  form  insoluble  basic  salts, 
of  a  lighter  blue  than  the  hydroxide. 

Ammonium  hydroxide  added  short  of  saturation  precipitates  the  pale 
blue  basic  salts;  added  just  to  saturation,  the  deep  blue  hydroxide  (in  both 
cases  like  the  fixed  alkalis);  added  to  supersaturation,  the  precipitate  dis- 
solves to  "an  intensely  deep  blue  solution  (separation  from  bismuth).  The 
blue  solution  is  a  cuprammonium  compound,  not  formed  unless  ammonium 
salts  be  present.  It  has  been  isolated  as  CuS04.(NH3)4  (§77,  56).  The  (loop 
blue  solution  probably  consists  of  this  compound  in  a  hydrated  condition, 
t.  e.  Cu(OH)2.2NH4OH.(NH4)2S04  ;  or  (NH4)4Cu(OH)4S04 .  Other  salts 
than  the  sulphate  form  the  corresponding  compounds:  CuCl2  +  4NH4OH 
=  Cu(OH),.2NH4OH.2NH4Cl .  The  blue  color  with  ammonium  hydroxide 
is  a  good  test  for  the  presence  of  copper  in  all  but  traces  (one  to  25,000), 
its  sensitiveness  is  diminished  by  the  presence  of  iron  (Wagner,  Z.,  1881, 
20,  351).  Ammonium  carbonate,  like  ammonium  hydroxide,  precipitates 
and  redissolves  to  a  blue  solution.  Carbonates  of  fixed  alkali  metals — as 
K2C03 — precipitate  the  greenish-blue,  basic  carbonate,  Cu2(OH)2C03 ,  of 
variable  composition,  according  to  conditions,  and  converted  by  boiling  to 
the  black,  basic  hydroxide  and  finally  to  the  black  oxide.  Barium  carbon- 
ate precipitates  completely,  on  boiling,  a  basic  carbonate. 

From  the  blue  ammoniacal  solutions  a  concentrated  solution  of  a  fixed 
alkali  precipitates  the  blue  hydroxide,  changed  on  boiling  to  the  black- 
oxide,  CuO  . 

6. — Oxalates,  cyanides,  f err o cyanides,  ferricyanides  and  thiocyanates  pre- 
cipitate their  respective  cuprous  salts  from  cuprous  solutions  not  too  strongly 
acid.  The  ferricyanide  is  brownish-red,  the  others  are  white.  The  thiocyamite 
is  used  to  separate  copper  from  palladium  (Wohler,  A.  Ch.,  1867,  (4),  10,  510); 
and  also  from  cadmium.  In  solutions  of  cupric  salts,  oxalates  precipitate  cupric 


§77,  60.  COPPER.  10? 

oxalate,  CuC2O4 ,  bluish-white,  insoluble  in  acetic  acid,  and  formed  from  mineral 
acid  salts  of  copper  by  oxalic  acid  added  with  alkali  acetates. 

Potassium  cyanide  forms  a  brownish  precipitate  of  cupric  cyanide, 
Cu(CN)2 ,  which  immediately  changes  to  the  yellowish  green  cupric  cuprous 
cyanide  with  the  evolution  of  cyanogen.  On  warming,  this  precipitate 
changes  to  the  white  cuprous  cyanide,  Cu2(CN)2,  with  further  evolution 
of  cyanogen.  Excess  of  potassium  cyanide  dissolves  the  precipitate  with 
the  formation  of  the  double  cuprous  salt,  K3Cu(CN)4?  or,  in  the  presence 
of  a  smaller  excess  of  potassium  cyanide,  K,Cu(CN)3.  (Kunschert,  Z. 
Anorg.,  41,  359,  1904).  The  potassium  cyanide  also  dissolves  cupric 
oxide,  hydroxide,  carbonate,  sulphide,  etc.,  changing  rapidly  to  the  double 
cuprous  cyanide  in  solution  in  the  alkali  cyanide.  Hydrogen  sulphide 
does  not  precipitate  the  copper  from  solutions  of  copper  salts  in  potassium 
cyanide  on  account  of  the  very  slight  concentration  of  the  copper  ions  in 
such  solutions  (5.10~20 .  Kunschert,  ibid.).  This  serves  as  a  separation 
from  Cadmium  which  is  precipitated  by  hydrogen  sulphide  from  a  potas- 
sium cyanide  solution.  Potassium  ferrocyanide  precipitates  cupric  ferro- 
cyanide,  Cu2Fe(CN)(! ,  reddish-brown,  insoluble  in  acids,  decomposed  by 
alkalis;  a  very  delicate  test  for  copper  (1  to  200,000)  ;  forming  in  highly 
dilute  solutions  a  reddish  coloration  (Wagner,  Z.,  1881,  20,  351). 
Potassium  ferricyanide  precipitates  cupric  ferricyanide,  Cu3(Fe(CN)6)2 , 
yellowish-green,  insoluble  in  hydrochloric  acid. 

Potassium  thiocyanate,  with  cupric  salts,  forms  a  mixed  precipitate  of 
cuprous  thiocyanate,  white,  and  a  black  precipitate  of  cupric  thiocyanate, 
which  gradually  changes  to  the  white  cuprous  compound,  soluble  in  NH4OH; 
in  the  presence  of  hypophosphorous  or  sulphurous  acid  the  cuprous  thiocyanate 
is  precipitated  at  once  (distinction  from  cadmium  and  zinc)  (Hutchinson,  J.  (7., 
1880,  38,  748).  Ammonium  benzoate  (10  per  cent  solution)  precipitates  copper 
salts  completely  from  solutions  slightly  acidified  (separation  from  cadmium) 
(Gucci,  B.,  1884,  IT,  2659). 

If  to  a  solution  of  cupric  salt  slightly  acidulated  with  hydrochloric  acid,  an 
excess  of  a  solution  of  nitroso-B-naphthol  in  50  per  cent  acetic  acid  be  added, 
the  copper  will  be  completely  precipitated  on  allowing  to  stand  a  short  time 
(separation  from  Pb  ,  Cd  ,  Hg  ,  Mn  ,  and  Zn)  (Knorre,  B.,  1887,  20,  283). 

Potassium  xanthate  gives  with  very  dilute  solutions  of  copper  salt  a  yellow 
coloration;  according  to  Wagner  (I.e.}  one  part  copper  in  900,000  parts  water 
may  be  detected. 

c. — Nitric  acid  rapidly  oxidizes  cuprous  salts  to  cupric  salts,  d. — A  solution 
of  cupric  sulphate  slightly  acidulated  with  hydrochloric  acid  is  precipitated  as 
cuprous  chloride  by  sodium  hypophosphite  (Cavazzi,  Gazzctta,  1886,  16,  167);  if 
the  slightly  acidulated  copper  salt  solution  be  boiled  with  an  excess  of  the 
hypophosphite  the  copper  is  completely  precipitated  as  the  metal.  Sodium 
phosphate,  Na_,HPO4  ,  gives  a  bluish-white  precipitate  of  copper  acid  phosphate, 
CuHPO4  ,  if  the  reagent  be  in  excess  and  Cu,(PO4)a  if  the  copper  salt  be  in 
excess.  Sodium  pyrophosphate  precipitates  cupric  salts,  but  not  if  tartrates 
or  thiosulphates  be  present  (separation  from  cadmium)  (Vortmann,  B.,  1888, 
21,  1103). 

e. — Cuprous  salts  (obtained  by  treating  cupric  salts  with  SnCL)  when  boiled 
with  precipitated  sulphur  deposit  the  copper  as  Cu.S  (separation  from  cad- 
mium) (Orlowski,  ./.  ('.,  1882,  42,  12:52).  Cuprous  salts  "are  precipitated  or  trans- 


108  COPPER.  §77,  6/. 

posed  by  hydrosulphuric  acid  or  soluble  sulphides,  forming  cuprous  sulphide,  * 
Cu2S ,  black,  possessing  the  same  solubilities  as  cupric  sulphide. 

With  cupric  salts  H2S  gives  CuS ,  black  (with  some  CiuS),  produced 
alike  in  acid  solutions  (distinction  from  iron,  manganese,  cobalt,  nickel) 
and  in  alkaline  solutions  (distinction  from  arsenic,  antimony,  tin). — Solu- 
tions containing  only  the  one-hundred-thousandth  of  copper  salt  are 
colored  brownish  by  the  reagent.  The  precipitate,  CuS ,  is  easily  soluble 
by  nitric  acid  (distinction  from  mercuric  sulphide)  ;  with  difficulty  soluble 
by  strong  hydrochloric  acid  (distinction  from  antimony)  ;  insoluble  in  hot 
dilute  sulphuric  acid  (distinction  from  cadmium)  ;  insoluble  in  fixed  alkali 
sulphides  and  ammonium  monosulphide,  but  slightly  soluble  in  ammonium 
polysulphide  and  fixed  alkali.  (Bossing,  Z.,  41,  1)  {  (distinction  from 
arsenic,  antimony,  tin)  ;  soluble  in  solution  of  potassium  c}ranide  (distinc- 
tion from  lead,  bismuth,  cadmium,  mercury). 

According  to  Noyes  (J.  Am.  Soc.  29,  170)  5  to  10  mg.  of  copper 
may  dissolve  in  ammonium  polysulphide  when  a  large  amount  of  copper 
is  present. 

Concerning  the  formation  of  a  colloidal  cupric  sulphide,  see  Spring  (B.,  1883, 
16,  1142).  According  to  Brauner  (C.  AT.,  1896,  74,  99)  cupric  salts  with  excess 
of  hydrogen  sulphide  always  yield  a  ver.y  appreciable  amount  of  cuprous  sul- 
phide. See  also  Ditte  (C.  r.,  1884,  98,  1492).  Solutions  of  cupric  salts  are 
reduced  to  cuprous  salts  by  boiling  with  sulphurous  acid  (Kohner,  C.  C.,  1886, 
813).  Sodium  thiosulphate  added  to  hot  solutions  of  copper  salts  gives  a  black 
precipitate  of  cuprous  sulphide.  In  solutions  acidulated  with  hydrochloric 
acid,  this  is  a  separation  from  cadmium  (Vortmann,  If.,  1888,  9,  165);  in  acetic 
acid  solution,  separation  from  cadmium  and  zinc  (Dovath,  Z.,  40,  141), 

/.— Hydrobromic  acid  added  to  cupric  solutions  and  concentrated  by 
evaporation  gives  a  rose-red  color.  Delicate  to  0.001  m.  g.  (Endemann 
and  Prochazka,  (7.  A7.,  1880,  42,  8).  Of  the  common  metals  only  iron 
interferes.  Potassium  bromide  and  sulphuric  acid  may  be  used  instead 
of  hydrobromic  acid. 

Hydriodic  acid  and  soluble  iodides  precipitate,  from  concentrated  solu- 
tions of  copper  salts,  cuprous  iodide,  Cul ,  white,  colored  dark  brown  by  the 
iodine  separated  in  the  reaction  J  (a).  The  iodine  dissolves  with  color  in 
excess  of  the  reagent,  or  dissolves  colorless  on  adding  ferrous  sulphate  or 
soluble  sulphites,  by  entering  into  combination.  Cuprous  iodide  dissolves 
in  thiosulphates  (with  combination). 

The  cuprous  iodide  is  precipitated,  free  from  iodine,  and  more  com- 

*  Freshly  precipitated  cuprous  sulphide  transposes  silver  nitrate  forming1  silver  sulphide, 
metallic  silver  und  cupric  nitrate ;  with  cupric  sulphide,  silver  sulphide  and  cupric  nitrate  are 
formed  (Schneider,  Pogg.,  1874, 152,  471).  Freshly  precipitated  sulphides  of  Fe  Co,  Zii,  Cd. 
Pb,  Bi,  Sn",  and  Snlv,  when  boiled  with  CuCl  in  presence  of  NaCl  give  Cu.,S  and  chloride  of 
the  metal:  with  CuCl2 ,  CuS  and  a  chloride  of  the  metal  are  formed,  except  that  SnS  gives 
Cu,S,  CuCl  and  S»IV  (Raschig,  B.,  1884,  17,  697). 

t  Thio  salts  having  the  formulas  (NH02  CiiaS?  and  NaiCujS?  are  formed  (Rossing,  Z.  anorg. 
25,  407.) 

t  The  precipitation  Is  incomplete  unless  the  free  iodine,  one  of  the  products  of  the  reaction,  is 
removed  by  means  of  a  reducing  agent  (§44). 


§77,  0.  COPPER.  109 

pletely,   bj   adding   reducing   agents   with    iodides;   as,   Na2SO.} ,   H0S03 , 

FeS04  (6). 

(a)     2CuS04  +  4KI  =  2CuI  +  I2  -f  2K2S04 

(6)     2CuS04  +  2KI  +  2FeSO,  —  2CuI  +  K2S04  +  l;e2(S04), 

2CuSO4  +  4KI  -f  H2S03  +  H20  =  2CuI  +  2K2SO4  +  H2S04  +  2HI 

ff. — Arsenites,  as  KAsO2  ,  or  arsenous  acid  with  just  sufficient  alkali  hydrox- 
ide to  neutralize  it,  precipitate  from  solutions  of  cupric  salts  (not  the  acetate) 
the  green  copper  arsenite,  chiefly  CuHAs03  (Scheele's  green,  "Paris  green"), 
readily  soluble  in  acids  and  in  ammonium  hydroxide,  decomposed  by  strong 
potassium  hydroxide  solution.  From  cupric  acetate,  arsenites  precipitate,  on 
boiling,  copper  aceto-arsenite,  (CuOAs,Og),Chl(C,H,Oj)a  ,  Schweinfurt  green  or 
Imperial  green,  "  Paris  green,"  dissolved  by  ammonium  hydroxide  and  by 
acids,  decomposed  by  fixed  alkalis. 

Soluble  arsenates  precipitate  from  solutions  of  cupric  salts  cupric  arsenate, 
bluish-green,  readily  soluble  in  acids  and  in  ammonium  hydroxide. 

*• — Potassium  bichromate  does  not  precipitate  solutions  of  cupric  salts: 
normal  potassium  chromate  forms  a  brownish*red  precipitate,  soluble  in  am- 
monium hydroxide  to  a  green  solution,  soluble  in  dilute  acids. 

7.  Ignition. — Ignition    with    sodium    carbonate    on    charcoal    leaves    metallic 
copper  in  finely  divided  grains.     The  particles  are  gathered  by  triturating  the 
charcoal  mass  in  a  small  mortar,  with  the  repeated  addition  and  decantation 
of  water  until  the  copper  subsides  clean.     It  is  recognized  by  its  color,  and 
its  softness  under  the  knife.     Copper  readily  dissolves,  from  its  compounds  in 
beads  of  borax  and  of  microco&mic  salt,  in  the  outer  flame  of  the  blow-pipe. 
The  beads  are  green  while  hot,  and   blue  when  cold.     In  the  inner  flame  the 
borax  bead  becomes  colorless  when  hot;  the  microcosmic  salt  turns  dark  green 
when  hot,  both  having  a  reddish-brown  tint  when  cold  (Cu20)   (helped  by  add- 
ing tin).     Compounds  of  copper,  heated  in  the  inner  flame,  color  the  outer  flame 
green.     Addition  of  hydrochloric  acid  increases  the  delicacy  of  the  reaction, 
giving  a  greenish-blue  color  to  the  flame. 

8.  Detection. — Copper  is  precipitated  from  its  solutions  by  H2S ,  form- 
ing CuS  .    By  its  insolubility  in  (NH4)2Sxand  solubility  in  hot  dilute  HN03 
it  is  separated  with  Pb  ,'Bi ,  and  Cd  from  the  remaining  metals  of  the  tin 
and  copper  group.     Dilute   H2S04  with   C2H5OH  removes   the  lead  and 
ammonium  hydroxide  precipitates  the  bismuth  as  Bi(OH)3 ,  leaving  the 
Cn  and  Cd  in  solution.     The  presence  of  the  Cu  is  indicated  by  the  blue 
color  of  the  ammoniacal  solution,  by  its  precipitation  as  the  brown  ferro- 
cyanide  after  acidulation  with  HC1  (66) ;  and  by  its  reduction  to  Cu°  with 
Fe°,  from  its  neutral  or  acidulated  solutions  (10).     Study  the  text  on 
reactions  (6)  and  §102  and  §103. 

9.  Estimation. — (1)  It  is  precipitated   on  platinum   by  the  electric   current 
and  weighed  as  the  metal,  or  by  means  of  zinc,  the  excess  of  zinc  being  dissolved 
by  dilute  hydrochloric  acid.      (2)   It  is  converted  into  CuO  and  weighed  after 
ignition,  or  the  oxide  is  reduced  to  the  metal  in  an  atmosphere  of  hydrogen  and 
weighed  as  such.     (3)  It  may  be  precipitated  either  by  H2S  or  Na2S2Os,   and, 
after  adding  free  sulphur  and  igniting  in  hydrogen  gas,  weighed  as  cuprous  sul- 
phide, or  it  may  be  precipitated  by  KCNS  in  presence  of  H?SOs  or  H  PO2  ,   and, 
after   adding  S  ,  ignited   in  H  and  weighed  as  Cu2S  .    Cu?O  ,  CuO  ,  Cu(NO3)2  , 
CuCO3  ,  CuSO4  and  many  other  cupric  salts,  are  converted  into  Ou2S  by  adding  S 
and  igniting  in   hydrogen  gas.      (4)   By  adding  KI  to  the  cupric  salt  and  titrating 
the  liberated  I  by  Na2S2O3  ;    not  permissible  with  acid  radicals  which  oxidize  HI  . 
(5)  By    precipitation    as    Cul    and    weighing  after   drying  at   150°    (Browning, 


110  CADMIUM.  §77,  10. 

Am.  S.,  1893  [3],  46,  280).  (6)  By  titrating-  in  concentrated  HBr  ,  using  a 
solution  of  SnClo  in  concentrated  HC1;  the  end  reaction  is  sharper  than  Avith 
SnCL,  alone  (Etard  and  Lebeau,  C.  r.,  1890,  110,  408).  (7)  By  titration  with 
Na,S.  Zinc  does  not  interfere  (Borntrager,  Z.  angeio.,  1893,  517).  (8)  By 
reduction  with  SO,  and  precipitation  with  excess  of  standard  NH4CNS;  dilu- 
tion to  definite  volume  and  titration  of  the  excess  of  NH4CNS  in  an  aliquot 
part,  with  AgNO3  (Volhard,  A.,  1878,  190,  51).  (9)  Small  amounts  are  treated 
with  an  excess  of  NH4OH  and  estimated  colorimetrically  by  comparing  with 
standard  tubes. 

10.  Oxidation.— Solutions  of  Cu"  and  Cu'  are  reduced  to  the  metallic 
state  by  Zn  ,  Cd ,  Sn  ,  Al ,  Pb  ,  Fe ,  Co  ,  Ni ,  Bi ,  Mg  *,  P  ,  and  in  presence 
of  KOH  by  K2Sn02 .  A  bright  strip  of  iron  in  solution  of  cupric  salts 
acidulated  with  hydrochloric  acid,  receives  a  bright  copper  coating,  recog- 
nizable from  solutions  in  120,000  parts  of  water.  With  a  zinc-platinum 
couple  the  copper  is  precipitated  on  the  platinum  and  its  presence  can  be 
confirmed  by  the  use  of  H2S04 ,  concentrated,  and  KBr ,  an  intense  violet 
color  is  obtained  (Creste,  J.  C.,  1877,  31,  803).  Cu"  is  reduced  to  Cu'  by 
Cu°  (Boettger,  J.  (7.,  1878,  34,  113),  by  SnCl2  in  presence  of  HC1 ,  in 
presence  of  KOH  by  As203  and  grape  sugar,  by  HI ,  and  by  S02 .  Metallic 
copper  is  oxidized  to  Cu"  by  solutions  of  Hg",  Hg',  Ag',  PtIV,  and  An'", 
these  salts  being  reduced  to  the  metallic  state.  Ferric  iron  is  reduced  to 
the  ferrous  condition  (Hunt,  Am.  $.,  1870,  99,  153).  Copper  is  also  oxi- 
dized by  many  acids. 

§78.  Cadmium.    Cd  —  112.4.     Valence  two. 

1.  Properties.  —  '-pecific  gr  vtiy,  liquid,  7.989;  cooled,    8.67;  hammered,  8.6944. 
Melting  point,  320.9°  (Cir.  B.  S.,  36,  1915).     Boiling  point,  763°  to  772°  (Car- 
nelley   and    Williams,    «/.    C.,    1878,    33,    284).     Specific   heat   is   0.0567.     Vapor 
density  (H  =  1),  55.8  (Deville  and  Troost,  A.  Ch.,  1860,   (3),  68,  257).     From 
these  data  the  gaseous  molecule  of  cadmium  is  seen  to  consist   of  one  atom 
(Richter,  Anonj.   Chan.,   1893,  363).     It  is  a  white  crystalline   metal,   soft,    but 
harder  than  tin  or  zinc;  more  tenacious  than  tin;  malleable  and  very  ductile, 
can  easily  be  rolled  out  into  foil  or  drawn  into  fine  wire,  but  at  SO0  it 'is  brittle. 
Upon  bending-  it  gives  the  same  creaking*  sound  as  tin.     It  may  be  completely 
distilled   in  a  current  of  hydrogen  above  800°,   forming  silver  white   crystals 
(Kanimerer,  B.,  1874,  7,  1724).     Only  slightly  tarnished  by  air  and  water  at 
ordinary  temperatures.     When  ignited  burns  to  CdO  .     When  heated  it  com- 
bines directly  with  Cl ,  Br  ,  I ,  P  ,  S  ,  Se  ,  and  Te  .     It  forms  many  useful  alloys 
having-  low  melting-points. 

2.  Occurrence. — Found  as  greenockite  (CdS)  in  Greenland,  Scotland  and  Penn- 
sylvania; also  to  the  extent  of  one  to  three  per  cent  in  many  zinc  ores. 

3.  Preparation. — Reduced  by  carbon  and  separated  from  zinc  (approximately) 
by  distillation,  the  cadmium  being  more  volatile.     It  may  be  reduced  by  fusion 
with  H  ,  CO  ,  or  coal  gas. 

4.  Oxide  and  Hydroxide. — Cadmium  forms  but   one  oxide,  CdO  ,   either  by 
burning  the  metal  in  air  or  by  ignition  of  the  hydroxide,  carbonate,  nitrate, 
oxalate,  etc.     It  is  a  brownish-yellow  powder,  absorbs  CO,  from  the  air,  becom- 
ing white    (Gmelin-Kraut,  3,  64).     The   Jiydroxule  -Cd(OH)2    is   formed   by   the 
action  of  the  fixed  alkalis  upon  the  soluble  cadmium  salts;  it  absorbs  CO2  from 
the  air. 

5.  Solubilities. — a. — Metal. — Cadmium    dissolves    slowly    in    hot,    moderately 
dilute  hydrochloric  or  sulphuric  acid  with  evolution  of  hydrogen;  much  more 

*Warren,  C,  N.,  1895,  71,  92, 


§78,6  CADMIUM.  Ill 

readily  in  nitric  acid  with  generation  of  nitrogen  oxides.  It  is  soluble  in 
ammonium  nitrate  without  evolution  of  gas;  cadmium  nitrate  and  ammonium 
nitrite  are  formed  (Morin,  C.  r.,  1885,  100,  1497).  ft. — The  oxide  and  Jiiidrn.ridc 
are  insoluble  in  water  and  the  fixed  alkalis,  soluble  in  ammonium  hydroxide, 
readily  soluble  in  acids  forming-  salts;  soluble  in  a  cold  mixture  of  fixed  alkali 
and  alkali  tartrate,  reprecipitated  upon  boiling  (distinction  from  copper) 
(Behal,  J.  Plianii.,  1885,  (5),  11,  553).  C. — Salts. — The  sulphide,  carbonate, 
oxalate,  phosphate,  cyanide,  ferrqcy_anide  and  ferricyanide  are  insoluble  (§27) 
in  Avater,  soluble  in  hydrochloric  and  nitric  acids,  and  soluble  in  NH4OH  , 
except  CdS  .  The  chloride  and  bromide  are  deliquescent,  the  iodide  is  perma- 
nent; they  are  soluble  in  water  and  alcohol. 

6.  Reactions,  a. — The  fixed  alkali  hydroxides— in  absence  of  tartaric 
and  citric  acids,  and  certain  other  organic  substances — precipitate,  from 
solutions  of  cadmium  salts,  cadmium  hydroxide,  Cd(OH)2 ,  white,  insoluble 
in  excess  of  the  reagents  (distinction  from  tin  and  zinc).  Ammonium 
hydroxide  forms  the  same  precipitate  which  dissolves  in  excess.  If  the 
concentrated  cadmium  salts  be  dissolved  in  excess  of  ammonium  hydroxide 
with  gentle  heat  and  the  solution  then  cooled,,  crystals  of  the  salt,  with 
variable  amounts  of  ammonia,  are  obtained;  e.  g.,  CdCl2(NH3)3 , 
CdS04(NH3)4,  Cd(N03)2(NH3)6  (Andre,  C.  r.,  1887,  104,  908  and  987; 
"Kwasnik,  Arch.  Pharm.,  1891,  229,  569).  The  fixed  alkali  carbonates  pre- 
cipitate cadmium  carbonate,  CdC03 ,  white,  insoluble  in  excess  of  the 
reagent,  ammonium  carbonate  forms  the  same  precipitate  dissolving  in 
excess.  Barium  carbonate,  in  the  cold,  completely^precipitates  cadmium 
salts  as  the  carbonate. 

6. — Oxalic  acid  and  oxalates  precipitate  cadmium  oxalate,  white,  soluble  in 
mineral  acids  and  ammonium  hydroxide.  Potassium  cyanide  precipitates 
cadmium  cyanide,  white,  soluble  in  excess  of  the  reagent  as  Cd(CN).,.2KCN; 
ffirrocyanides  form  a  white  precipitate;  ferricyanides  a  yellow  precipitate, 
both  soluble  in  hydrochloric  acid,  and  in  ammonium  hydroxide.  Potassium 
sulphocyanate  does  not  precipitate  cadmium  salts  (distinction  from  copper). 
Cadmium  salts  in  presence  of  tartaric  acid  are  not  precipitated  by  fixed  alkali 
hydroxides  in  the  cold;  on  boiling,  cadmium  oxide  is  precipitated  (separation 
from  copper  and  zinc)  (Aubel  and  Ramdohr,  A.  Cli.,  1858,  (3),  52,  109). 
c. — Nitric  acid  dissolves  all  the  known  compounds  of  cadmium,  d. — Soluble 
phosphates  precipitate  cadmium  phosphate,  white,  readily  soluble  in  acids. 
Sodium  pyrophosphate  precipitates  cadmium  salts,  soluble  in  excess  and  in 
mineral  acids,  not  in  dilute  acetic.  The  reaction  is  not  hindered  by  the  pres- 
ence of  tartrates  or  of  thiosulphates  (separation  from  Cu)  (Vortmann,  B.,  1888, 
21,  1104). 

e. — Hydrogen  sulphide  and  soluble  sulphides  precipitate,  from  solutions 
neutral,  alkaline,  or  not  too  strongly  acid,  cadmium  sulphide,  yellow; 
insoluble  in  excess  of  the  precipitant  (Fresenius,  Z.,  1881,  20,  236),  in 
ammonium  hydroxide,  or  in  cyanides  (distinction  from  copper);  soluble  in 
hot  dilute  sulphuric  acid  and  in  a  saturated  solution  of  sodium  chloride  * 
(distinction  from  copper)  (Cushman,  Am.,  1896,  17,  379). 

*  Owing  to  the  formation  of  incompletely-dissociated  CdCl3.  CdI2  is  still  less  dissociated 
and  accordingly  CclS  dissolves  more  readily  in  HI  than  in  HC1  and  much  more  readily  than  in 
HNO3  of  the  same  concentration.  On  the  other  hand,  of  course,  precipitation  of  the  sulphide 
takes  place  with  more  difficulty  from  the  iodide  than  from  the  other  salts. 


112  CADMIUM.  §78,  6/. 

Sodium  thiosulphate,  Na2SsO8  ,  does  not  precipitate  solutions  of  cadmiuir 
salts  (Follenius,  Z.,  1874,  13,  438),  but  in  excess  of  this  reagent,  ammonium 
salts  being-  absent,  sodium  carbonate  completely  precipitates  the  cadmium  as 
carbonate  (distinction  from  copper)  (Wells,  C.  N.,  1891,  64,  294).  Cadmium 
salts  with  excess  of  sodium  thiosulphate  are  not  precipitated  upon  boiling1 
with  hydrochloric  acid  (distinction  from  copper)  (Orlowski,  J.  C.,  1882,  42,  1232). 
f> — The  non-precipitation  by  iodides  is  a  distinction  from  copper,  g. — Soluble 
arsenites  and  arsenates  precipitate  the  corresponding  cadmium  salts,  readily 
soluble  in  acids  and  in  ammonium  hydroxide,  h. — Alkali  chromates  precipitate 
yellow  cadmium  chromate  from  concentrated  solutions  only,  and  soluble  oiv 
addition  of  water. 

*• — A  solution  of  copper  and  cadmium  salts,  very  dilute,  when  allowed  to 
spread  upon  a  filter  paper  or  porous  porcelain  ple.te,  gives  a  ring  of  the  cad- 
mium salt  beyond  that  of  the  copper  salt,  easily  detected  by  hydrogen  sulphide 
(Bag-ley,  J.  C.,  1878,  33,  304). 

7.  Ignition. — On  charcoal,  with  sodium  carbonate,  cadmium  salts  are  reduced 
by  the  blow-pipe  to  the  metal,  which  is  usually  vaporized  and  reoxidized  nearly 
as  fast  as  reduced,  thereby  forming  a  characteristic  brown  incrustation  (CdO). 
This  is  volatile  by  reduction  only,  being  driven  with  the  reducing  flame.     Cad- 
mium  oxide   colors   the   borax   bead   yellowish   while   hot,    colorless   when   cold; 
microcosmic  salt,  the  same.     If  fused  with  a  bead  of  K2S  ,  a  yellow  precipitate 
of  CdS  is  obtained  (distinction  from  zinc)  (Chapman,  J.  C.,  1877,  31,  490). 

8.  Detection. — Cadmium  is  precipitated  from  its  solutions  by  H2S  forming 
CdS.  By  its  insolubility  in  (NH4)2S  or  (NH4)2SZ    and  solubility  in  hot  dilute 
HN03  it  is  separated  with  Pb ,  Bi ,  and  Cu  from  the  other  metals  of  the 
second  group.     Dilute  H2S04  with  C2H5OH  removes  the  lead  and  NH4OH 
precipitates  the  bismuth  as  Bi(OH)3 ,  leaving  the  Cu  and  Cd  in  solution. 
If  copper  be  present,  KCN  is  added  until  the  solution  becomes  colorless, 
when  the  Cd  is  detected  by  the  formation  of  the  yellow  CdS  with  H2S  . 
If  Cu  be  absent  the  yellow  CdS  is  obtained  at  once  from  the  ammoniacal 
solution  with  H2S  .     See  also  Gi. 

9.  Estimation. — (1)   It  is  converted  into,  and  after  ignition  weighed  as    an 
oxide.     (2)  Converted  into,  and  after  drying  at  100°,  weighed  as  CdS.     (3)  Pre- 
cipitated as  CdC2O4  and  titrated  by  KMnO4.    (4)  Electrolytically  from  a  slightly 
ammoniacal  solution  of  the  sulphate  or  from  the  oxalate  rendered  acid  with 
oxalic  acid.     (5)  Separated  from  copper  by  KI;  the  I  removed  by  heating;  the 
excess   of   KI    removed    by    KNO2    and    H2SO4;    the   cadmium   precipitated    by 
Na2CO8  and  ignited  to  CdO  (Browning,  Am.  8.,  1893,  146,  280).     (6)  By  adding 
a  slight  excess  of  H2SO4  to  the  oxide  or  salt,  and  evaporation  first  on  the  water 
bath  and  then  on  the  sand  bath,  weighed  as  CdS04  (Follenius,  Z.,  1874,  13,  277). 

10.  Oxidation. —Metallic  cadmium  precipitates  the  free  metals  from 
solutions  of  Au ,  Pt ,  Ag ,  Hg ,  Bi ,  Cu ,  Pb ,  Sn ,  and  Co  ;  and  is  itself 
reduced  by  Zn ,  Mg ,  and  Al . 


,  1.         PRECIPITATION  OF  METALS   OF  SECOND   GROUP. 


113 


§79.  Comparison  of  Certain  Reactions  of  Bismuth,  Copper,  and  Cadmium. 

Taken  in  Sohilio-ns  of  their  Chlorides,  Nitrates,  Sulphates,  or  Acetates. 


Bi 

Cu 

Cd 

KOH  .  r  NaOH,  in 

Bi(OH)8,  white. 

Cu(OH);i,   dark 
blue 

Cd(OH)a,  white. 


NH4OH,  in  excess 
Dilution    of    satu- 
rated solutions 

Bi(OH)3,  white. 
BiOCl,  white  (§76, 

:,'/) 

Blue  solution. 

Colorless  solution. 

Iodides    

Partial     precipita- 

Precipitation       of 

tion  in  solutions 

Cul   with  libera- 

Sulphides   
Iron  or  zinc  

not  very  strongly 
acid  (§76,  6f). 
Bi.2S3,     black,     in- 
soluble in  KCN. 

Bi   spongy  precipi- 

tion     of     iodine 
(§77,  6/). 
Cu2S       and       CuS, 
black,  soluble  in 
KCN. 
Cu    bright  coating 

CdS,  yellow,  insol- 
uble in  KCN. 

Cd      gray    sponge 

Glucose,  KOH,  and 

Vtonf 

tate. 
Bi,  black. 

(§77,  10). 
Cu20,  yellow  (§77, 

K.h\ 

with  zinc,  no  ac- 
tion with  iron. 

K28nO,   -f  KOH.. 

Bi,  black. 

50). 
Cu,    precipitated 

SYSTEMATIC  ANALYSIS  OF  THE  METALS  OF  THE  TIN  AND  COPPER  GROUT. 

The  precipitation  of  the  metals  of  the  second  group  (Tin  and  Copper 
Group)  by  hydrosulphuric  acid,  and  their  separation  into  Division  A 
(Copper  Group)  and  Division  B  (Tin  Group).  See  §312. 

§80.  Manipulation. — The  filtrate  from  Group  1  (§62),  or  the  original 
solution,  if  the  metals  of  the  silver  group  be  absent  is  warmed  nearly  to 
boiling  and  saturated  with  H2S  gas.  The  volume  of  the  solution  should 
be  about  50  c.c.  and  it  should  contain  about  6  per  cent  of  concentrated 
HC1  by  volume.  After  passing  H2S  for  about  15  minutes  through  the  hot 
solution,  allow  it  to  cool,  dilute  with  an  equal  volume  of  water  and  again 
pass  H2S  for  some  time  through  the  cold  solution.  Shake  well  to  coagulate 
the  precipitate  and  filter.  Pass  H2S  again  through  the  filtrate,  filter  and 
repeat  until  no  further  H2S  precipitate  is  obtained. 

2H3AsO4  +  xHCl  +  5H,S  =  As2S6  +  xHCl  +  8H2O 
or          2H3AsO4  +  xHCl  +  5H2S  =  As2S,  +  xHCl  +  S,  +  8HaO 
SnCl4  +  2H2S  =  SnS2  -f-  4HC1 
SnCl2  +  H2S  -  SnS  +  2HC1 
2Bi(NO3)3  +  3H2S  =  Bi2S3  +  6HNO, 
CdSO,  +  H2S  =  CdS  +  H,SO4 

§81.  Notes. — /.  Hydrosulphuric  acid  gas  should  be  used  in  precipitating  the 
metals  of  the  second  group.  It  may  be  generated  in  a  Kipp  apparatus,  using 
ferrous  sulphide,  FeS  ,  and  dilute  commercial  sulphuric  acid  (1-12).  Com- 
mercial hydrochloric  acid  may  be  used  instead  of  sulphuric.  The  gas  should 


114  PRECIPITATION  OF  METALS   OF  SECOND   GROUP.         §81,  #. 

be  passed  through  a  wash  bottle  containing  water  to  remove  any  acid  that  may 
be  carried  over  mechanically.  It  should  always  be  conducted  through  a  capil- 
lary tube  into  the  solution  to  be  analyzed  contained  in  a  flask.  Less  gas  is 
required  and  the  solution  is  less  liable  to  be  thrown  from  the  test  tube  by  the 
excess  of  unabsorbed  gas. 

2.  In  treating  the  unknown  solution  with  H2S,  it  should  be  passed  into  the 
liquid  until,  upon  shaking  the  flask,  capped  with  the  thumb,  there  is  no  forma- 
tion of  a  partial  vacuum  due  to  the  further  absorption  of  the  gas  by  the  liquid. 

8.  H2S  is  decomposed  by  HNO3  or  HNO3  +  HC1  (nitrohydrochloric  acid) 
(§257,  6B),  hence  these  acids  must  not  be  present  in  excess.  If  these  acids 
were  used  in  preparing  the  solutions  for  analysis,  they  must  be  removed  by 
evaporation.  Sulphuric  acidulation  is  not  objectionable  to  precipitation  with 
H2S  ,  but  could  not  be  used  until  absence  of  the  metals  of  the  calcium  group 
(Group  V.)  has  been  assured.  If  this  group  is  present  strontium  and  especially, 
barium,  will  invariably  be  present  in  the  H2S  precipitate  on  account  of  the  oxi- 
dation of  the  sulphur  to  sulphuric  acid.  For  this  reason,  oxidizing  agents  must 
be  removed  from  the  solution  so  far  as  possible.  If  ferric  chloride  is  present,  15 
milligrams  of  barium  may  be  present  in  this  precipitate  as  sulphate.  Curt  man  and 
Frankel  (J.  Am.  Soc.,  33,  724,  1911.)  For  detection  of  the  barium  see  §301,  5. 

4-  The  precipitation  of  the  silver  group  has  left  the  solution  acid  with  HC1 
and  prepares  the  solution  for  precipitation  with  H2S  if  other  acids  are  not 
present  in  excess.  A  moderate  excess  of  HC1  is  necessary  to  insure  the  precipi- 
tation of  arsenic  if  present  in  the  arsenic  condition.  For  this  purpose  the 
solution  must  be  hot  and  must  contain  at  least  6  per  cent  by  volume  of  con- 
centrated HC1.  Under  these  conditions  the  arsenic  precipitates  slowly  (69). 

The  strong  acid,  especially  when  hot,  prevents  the  precipitation  of  the  other 
metals,  especially  tin,  lead  and  cadmium.  For  this  reason,  the  solution  must  be 
cooled  and  diluted  and  again  saturated  with  H2S  in  order  to  precipitate  these 
metals.  The  solution  must  not  be  too  largely  diluted  or  traces  of  Co  ,  Ni  and  Zn 
will  be  precipitated.  About  one  part  of  HC1  to  25  of  the  solution  must  be  present 
to  prevent  the  precipitation  of  Zn  ,  and  it  is  seldom  advisable  to  use  more  than 
one  part  of  HC1  to  ten  of  the  solution  *  (this  refers  to  the  reagent  HC1  ,  §324). 

5.  The    precipitated   sulphides   of   the    metals   of   the    tin   and    copper   group 
(second  group)  present  a  variety  of  colors,  which  aid  materially  in  the  further 
analysis  of  the  group.     CdS,    SnS2,  As2S3 ,   and  As2S5  are  lemon-yellow;  Sb-S5 
and  SboS?   are  orange;    SnS  ,   HgS ,  PbS,   Bi2S3  ,   Cu2S  and  CuS   are   black  to 
brownish-black.     If  too  much  HC1  be  present,  lead  salts  frequently  precipitate  a 
red  double  salt  of  lead  chloride  and  lead  sulphide  (§67,  6e}.     Mercuric  chloride  at 
first  forms  a  white  precipitate  of  HgCl2.2HgS,  changing  from  yellow  to  red,   and 
finally  to  black  with  more  H2S  ,  due  to  the  gradual  conversion  to  HgS  (§58,  6e). 

6.  Addition  of  water  to  the  solution  before  passing  in  H2S   may  cause    the 
precipitation  of  the  oxychlorides  of  Sb  ,  Sn  or  Bi  (5d;    §70,  §71  and  §76).     These 
should  not  be  redissolved  by  the  addition  of  more  HC1  ,   as  they  are    readily 
transposed  to  the  corresponding  sulphides  by  H2S  ,  and  the  excess  of  acid  nec- 
essary to  their  resolution  may  prevent  the  precipitation  of  cadmium  or  cause 
the  formation  of  the  red  precipitate  with  lead  chloride. 

7.  The  presence  of  a  strong  oxidizing  agent  as  HNO3  ,  KL>Cr2O7  ,  FeCl3  ,    etc., 
causes  with  H2S   the  formation   of  a  white   precipitate   of  sulphur    (§125,    6e), 
which  is  often  mistaken  as  indicating  the  presence  of  a  second  group  metal. 

*  Addition  of  a  strong  acid,  containing  H  ions  in  large  quantity,  diminishes  the  already  slight 
(Dissociation  of  the  H2S  (§44),  thus  decreasing  in  number  the  S  ions,  whose  concentration  multi- 
plied by  that  of  the  metal  ions  must  equal  the  solubility-product  of  the  sulphide  in  question 
before  precipitation  can  take  place.  Precipitation  of  some  of  the  sulphides  of  the  Tin  and 
Copper  Group  may  be  entirely  prevented  in  this  way. 

It  frequently  happens  that  addition  of  water  alone  will  cause  precipitation  of  these  sulphidos 
from  a  strongly  acid  solution  which  has  been  saturated  with  H2S.  This  appears  strange  in  view 
of  the  fact  that  the  acid  which  prevented  precipitation  and  the  acid  which  finally  produced  it 
were  both  diluted  by  the  added  water  in  the  same  proportion.  Hut  as  a  matter  of  fact  dilution 
does  not  have  the  same  effect  on  a  strong  acid  as  on  a  weak  one.  Dissociation  is  always  in-- 
creased by  dilution,  but  in  much  greater  ratio  in  the  case  of  a  weakly-dissociated  body  as  H2S 
than  where  the  dissociation  of  the  substance  is  already  practically  complete,  as  in  the  case  of 
the  strong  acid.  Dilution  in  the  case  mentioned  increases  the  relative  concentration  of  the  S 
ions  and  so  the  solubility-product  is  reached  and  precipitation  results. 


§83,2.          PRECIPITATION  OF  METALS   OF  SECOND   GROUP.  115 

If  the  original  solution  be  dark  colored,  it  is  advisable  to  warm  with  hydro- 
chloric acid  and  alcohol  (§126,  6/  and  10)  to  effect  reduction  of  a  possible  higher 
oxidized  form  of  Cr  or  Mn  before  the  precipitation  with  H2S  ,  thus  avoiding 
the  unnecessary  precipitation  of  sulphur. 

8.  Complete  precipitation  of  the  metals  of  the  second  group  with  H2S  may 
fail:  (1)  from  incomplete  saturation  with  the  gas  (§81,  2);  (2)  from  the  pres- 
ence of  too  much  HC1  (§81,  4);  (3)  from  the  presence  of  much  pentad  arsenic 
(§69,  6e).  The  first  cause  of  error  may  be  avoided  by  careful  observance  of  the 
directions  in  note  (2).  Too  much  acid  may  be  present  because  excess  of  acid 
had  been  used  in  dissolving  the  unknown.  After  precipitating  the  first  group, 
excess  of  nitric  should  be  removed  by  evaporating  the  solution  nearly  to  dryness 
then  diluting  and  adding  the  required  amount  of  HC1.  As  a  further  precaution 
a  portion  of  the  filtrate  from  the  H2S  precipitate  should  be  diluted  with  several 
volumes  of  water  and  H^S  passed.  If  a  precipitate  is  obtained,  the  entire  solu- 
tion should  be  diluted  and  saturated  with  H2S.  Asv  must  be  precipitated  by 
passing  H2S  rapidly  through  the  hot  moderately  acid  solution  before  dilution  as 
long  as  the  slow  formation  of  the  arsenic  precipitate  continues. 

§82.  Manipulation. — -After  the  precipitate  has  been  well  washed  with 
hot  water  the  point  of  the  filter  is  pierced  with  a  small  stirring  rod  and 
the  precipitate  washed  into  a  beaker,  using  as  small  an  amount  of  water 
as  possible.  If  As,  Sb  and  Sn  are  present,*  ammonium  sulphide  (NH4)2S 
(§38,  2)  is  then  added  and  the  precipitate  digested  for  several  minutes 
with  warming: 

As2S8  +  2(NH4)2S2  =  (NH4)4As2S5  +  S2 
SnS  +  (NH4)2S2  =  (NH4),SnS3 
2SnS2  +  2(NH4)2S,  =  2(NH4)2SnS3  +  S2 
2Sb2S3  +  6(NH4)2S2  =:4(NH4)3SbS4  +  S2 
2MoS3  +  2(NH4)2S3  —  2(NHJ2MoS4  +  S2 

The  precipitate  is  then  filtered  and  washed  once  or  twice  with  a  small 
amount  of  (NH4)2S,  and  then  with  hot  water.  The  filtrate  consisting  of 
solutions  of  the  sulphides  of  As,  Sb ,  Sn ,  Au,  Ft,  Mo  (Gr,  Ir,  Se,  Te, 
IV,  !'),  constitutes  the  Tin  Group  (Division  A  of  the  second  group).  The 
precipitate  remaining  upon  the  filter,  consisting  of  the  sulphides  of  Hg , 
Pb,  Bi,  Cu,  Cd  (Os,  Pd,  Eli,  and  Ru),  constitutes  the  Copper  Group 
(Division  B  of  the  second  group,,  §95). 

§83.  IVo/rs. —  1.  The  precipitate  of  the  sulphides  of  the  tin  and  copper  group 
must  be  thoroughly  washed  with  hot  water  (preferably  containing1  H2S  and 
about  one  per  cent  of  reagent  HC1  to  prevent  the  formation  of  soluble  colloidal 
sulphides  (§69,  5c),  to  insure  the  removal  of  the  metals  of  the  iron  and  zinc 
groups,  which  would  be  precipitated  on  the  addition  of  the  ammonium  sulphide 
(§144). 

:.'.  Yellow  ammonium  sulphide,  (NH4)2SX,  forms  upon  allowing-  Iho  normal 
sulphide,  (NH4)2S  ,  to  stand  for  sometime,  or  it  may  be  prepared  for  imme- 
diate use  by  adding  sulphur  to  the  freshly  prepared  normal  sulphide  (§257,  4). 

*  This  operation  is  necessary  only  when  both  divisions  of  the  group  are  present,  and  is  to  be 
avoided  when  unnecessary.  Hence  a  little  of  the  2nd  group  precipitate  is  tested  by  warming 
with  1  or  2  cc.  (NH-OaSx  .  If  it  all  dissolves,  only  As,  Sb,  Sn  can  be  present;  if  nothing  dissolves, 
none  of  these  can  be  present;  if  part  dissolves,  then  the  whole  2nd  group  precipitate  must  be  so 
treated.  To  see  if  anything  has  dissolved  in  the  (NH4)2Sx  it  is  acidified  slightly  with  HC1  (test 
with  litmus);  a  milky,  white  precipitate  of  S  will  always  be  formed,  but  if  any  sulphides  are 
present  they  will  appear  as  a  flocculent,  colored  precipitate.  If  the  whole  2nd  group  precipitate 
s  treated  with  (NHOzSx  ,  the  solution  is  filtered  and  acidified  just  as  the  test  portion  was. 


116 


TABLE  FOR  THE  ANALYSIS  OF  THE  TIN  GROUP. 


>84. 


Pi 


».g 

3 


PH 


W      M 


p 


g»°* 

K    *"  s^,3 


I  -^ 
«£  S 

^    .  5 


0  »*« 

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TABLE  FOR  THE  ANALYSIS  OF  THE  TIN  GROUP. 


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118  DIRECTIONS  FOR   ANALYSIS    WITH  NOTES.  §83,  3. 

For  arsenic  sulphides  the  normal  ammonium  sulphide  may  be  employed,  but 
the  sulphides  of  antimony  are  soluble  with  difficulty,  and  stannous  sulphide  is 
scarcely  at  all  soluble  in  that  reagent;  while  they  are  all  readily  soluble  in  the 
yellow  polysulphide  (6e;  §69,  §70  and  §71). 

3.  Cupric  sulphide,  CuS  ,  is  sparingly  soluble  in  the  yellow  ammonium  sul- 
phide and  will   give  a  grayish-black  precipitate  upon   acidulation  with   HC1  . 
The  sulphides  of  the  tin  group  are  soluble  in  the  fixed  alkali   sulphides,  K,S 
and  Na2S;  cupric  sulphide  is  insoluble  in  these  sulphides.     Mercuric  sulphide, 
however,  is  much  more  soluble  in  fixed  alkali  sulphides  than  cupric  sulphide  is 
in  the   (N"H4),SX  .     If  copper  be  present  arid  mercury  be  absent,  it  is  recom- 
mended  to   use  KL,S  or  Na2S  instead  of    (NH4)2SX  for  the   separation   of   the 
second  group  of  sulphides  into  divisions  A  (tin  group)  and  B  (copper  group). 
But  if  Hg"   be  present,   the    (NH4)2SX   should   be   used,   and  the   presence   or 
absence  of  traces  of  copper  be  determined  from  a  portion  of  the  filtrate  from 
the  silver  group  before  the  addition  of  H2S  (§103). 

4.  The  sulphides  dissolve  more  readily  in  the  (NH4)2SX  when  the  solution  is 
warmed.     An  excess  of  the  reagent  is  to  be  avoided,  as  the  acidulation  of  the 
solution  causes  the  precipitation  of  sulphur  (§256,  3),  which  may  obscure  the 
precipitates  of  the  sulphides  present. 

§85.  Manipulation.  —  The  solution  of  the  sulphides  in  (NH4)2SX-  is  care- 
fully acidulated  with  hydrochloric  acid: 

2(NH4)2S2  +  4HC1  =  4NH4C1  +  S2  +  2ELS 

(NH4)4As2S3  +  4HC1  =  As2S,  +  4NH4C1  +  2H,S 
2(NH4)3SbS4  +  6HC1  =  Sb2S5  +  6NH4C1  +  3H2S 
(NH4)2SnS3  +  2HC1  =  SnS2  +  2NH4C1  +  H2S 

The  precipitate  obtained  when  the  metals  of  the  tin  group  are  present, 
is  usually  yellovr  or  orange-yellow  and  is  easily  distinguished  from  a  pre- 
cipitate of  sulphur  alone  (SnS  and  MoS3  are  brownish-black).  It  should 
be  well  washed  vita  hot  water  and  then  dissolved  in  hot  HC1  using  small 
fragments  o2  L£C103  (§69,  6e)  to  aid  in  the  solution: 

2As2S3  -:-  10C12  +  16H,,0  =  4H3As04  +  20HC1  +  3S2 
SnS2  +  4HC1  =  SnCl4  + 

rts2  -:-  ^ci2 


The  solution  iz  boiled  (to  insure  removal  of  the  chlorine  (§69,  10)  until  it 
no  longer  bleaches  litmus  paper. 

§86.  Notes.  —  1.  If  the  precipitate  obtained  is  white,  it  probably  consists  of 
sulphur  alone  and  indicates  absence  of  more  than  traces  of  the  metals  belong- 
ing to  this  group  (GeS2  is  white,  §111,  6). 

2.  Care  should  be  taken  not  to  use  too  much  HC1  in  precipitating  the  sul- 
phides from  the  (NH4)2SX  solution,  as  some  of  the  sulphides  (especially  SnS,,) 
are  quite  soluble  in  concentrated  HC1  . 

.?.  It  will  be  noticed  (§85)  that  the  lower  sulphides  of  Sb  and  Sn  are  oxidized 
by  the  (NH4)2SX  ,  and  are  precipitated  by  the  HC1  as  the  higher  sulphides 
Sb2S5  and  SnS2  respectively.  This  fact  may  be  most  readily  observed  l>\  the 
precipitation  of  a  solution  of  SnCL  with  H2S  ,  giving  a  brown  precipitate  of 
SnS  ,  then  dissolving  this  precipitate  in  (NH4)L.SX  and  reprecipitating  with  HC1 
as  the  orange-colored  SnS,.  . 

//.  Hot  reagent  HC1  (§324)  dissolves  the  sulphides  of  tin  quite  readily 
without  reduction;  the  sulphides  of  antimony,  slowly  forming  SbCl3  only;  and 
the  sulphides  of  arsenic  practically  not  at  all,  or  at  most  only  traces.  The 
sulphides  of  Au  and  Pt  are  not  soluble  in  HC1  .  MoS3  is  soluble  in  hot  con- 


DIRECTIONS  FOR   ANALYSIS   WITH  NOTES. 


119 


eentrated  HC1 .  The  relative  solubility  of  these  sulphides  in  HC1  is  used  as  the 
basis  of  the  following  separation  of  As  from  Sb  and  Sn  (§69,  6e,  also  bottom  of 
next  note,  5). 


Free. :   As2S5 ,  SboS5 ,  SnS2 . 

(sp.  gr.  1.2).     Expel  all  H2S  . 


Heat  for  a  few  moments  with  concentrated  HC1 
Dilute  a  little  and  filter. 


Residue  :—As2Sft ,  (S2).     Apply  either  of 
the  following  tests: 

(1)  Pour  warm  NH4OH  over  the  precipi- 
tate.     Add  H2O2  to  this  solution  and 
boil.     Cool  and  add  a  few  cubic  centi- 
meters of  NH4C1  and  a  little  MgCl2  and 
obtain  a  white  crystalline  precipitate  of 
MgNH4AsO4. 

(2)  Dissolve   As2S5   in  HC1  +  crystal   of 
KC1O3.      Boil.       Make    alkaline  with 
NH4OH  ,   and  add  NH4C1  and  MgCl2 
as  in  (1). 


Filtrate:— SbCl3 ,  SnCl4  .  Boil  to  be 
sure  of  complete  expulsion  of  H2S . 

Test  for  Sb: — Put  drop  of  solution  on 
silver  coin.  Bend  piece  of  tin  in  form 

of   i 1.      Touch  one    end    to    drop 

and  other  to  coin  outside  the  drop. 
Allow  to  stand  for  a  few  moments. 
Brown  or  black  spot  on  coin  is  due  to 
metallic  Sb . 

Test  for  Sn: — Heat  solution  with  iron 
wire  until  reduction  is  complete.  Filter 
and  to  filtrate  add  HgCL  A  white 
precipitate  of  HgCl  or  a  gray  precip- 
itate of  Hg  shows  tin. 


The  precipitated  sulphides  of  As  ,  Sb  ,  Sn  are  well  washed  with  hot  water  and 
removed  from  the  filter  ^to  a  casserole  by  a  spatula,  or,  if  the  amount  is  small, 
treated  with  the  filter;  a  convenient  amount  of  concentrated  HC1  (sp.  gr.  1.2)  is 
added  and  boiled  a  minute  or  two  to  expel  H2S  .  The  sulphides  of  Sb  and  Sn  are 
dissolved  to  form  the  chlorides  SbCl3  and  SnCl4  while  the  As2S3  is  hardly  attacked. 
Since  the  strong  acid  attacks  the  filter  the  solution  is  diluted  a  little,  which  should 
cause  no  reprecipitation  if  all  H2S  was  expelled,  filtered,  and  the  residue  well 
washed.  It  may  be  either  As2S3  and  S  ,  or  S  alone.  A  few  cc.  of  warm  NH4OH 
are  poured  over  it,  the  solution  being  passed  through  again  if  necessary.  The 
As2S3  dissolves  and  the  S  remains.  To  the  solution,  which  must  be  clear,  add  1  or  2 
cc.  H2O2  ,  2  to  3  cc.  NH4C1  ,  and  2  to  3  cc.  "magnesia  mixture,"  which  is  MgCl2  -f- 
NH4C1  4-  NH4OH  .  Cool,  and  let  stand  for  a  time.  The  Asv  is  precipitated 
as  NH.;MgAsO4 ,  a  white,  crystalline  precipitate  exactly  like  NH,MgPO4  in  ap- 
pearance. 

As2S6  +  16NH4OH  +  20H2O2  =  2(NH4)3AsO4  +  5(NH4)2SO4  +  28H2O . 
(NH4)3As04  +  MgCl2  +  [NH4OH  +  NH4C1]  =  MgNH4AsO4  +  2NH4C1 . 

The  filtrate  from  As2S3  is  to  be  tested  for  Sb  and  Sn  .  For  the  Sb  ,  place  a  few 
drops  on  a  clean  silver  coin;  it  should  produce  no  discoloration.  A  piece  of  tin, 
bent  into  the  shape  of  a  broad  U,  is  now  placed  on  the  coin  so  that  one  end  is  in 
the  center  of  the  drop  and  the  other  in  contact  with  the  silver  outside.  Allow 
to  stand  about  5  minutes.  If  Sb  is  present  it  will  be  deposited  as  a  brown  spot 
on  the  silver  covered  by  the  drop,  the  Sn  and  Ag  acting  as  a  galvanic  couple  to 
reduce  the  Sb  *  *  *  to  metal.  Another  test  consists  in  treating  the  solution  with 
pure,  fine  Fe  wire,  the  Sb  being  precipitated  in  black  metallic  form,  while  the 
Sn  *  '  *  *  is  merely  reduced  to  Sn  •  •  but  not  precipitated. 

Test  the  rest  of  the  solution  for  Sn  by  heating  with  fine  Fe  wire  until  the  solution 
is  colorless  or  greenish,  with  no  trace  of  yellow,  to  make  sure  that  all  the  Sn  •  •  •  '  is 
reduced  to  Sn  *  •  .  Ten  minutes  or  more  may  be  required.  Filter  and  add  the 
filtrate  slowly  (a  few  drops  at  a  time),  to  a  few  cc.  of  ammonium  molybdate, 
(NH4)2MoO4  solution.  A  deep  blue  color  or  precipitate  will  appear  if  Sn  *  •  is 
present,  due  to  the  reduction  of  the  MoO3  to  a  lower  exide.  Or,  instead  of  adding 
this  filtrate  to  molybdate  solution,  it  may  be  treated  with  HgCl2  ,  a  white  precip- 
itate of  HgCl  being  formed  if  Sn  *  •  is  present.  Note  that  this  is  reversing  the  test 
for  Hg  •  •  with  SnCl2 .  The  HgCl2  test  is  most  characteristic. 


120 


DIRECTIONS  FOR  ANALYSIS    WITH  NOTES. 


\,5. 


The  precipitation  of  Ag2S3  ,  unlike  that  of  the  other  sulphides,  is  not  prevented  by 
the  presence  of  any  amount  of  HC1 ,  however  large,  but,  on  the  contrary,  is  aided. 
It  may,  therefore,  be  necessary,  after  removing  all  other  sulphites  in  the  N/5  HC1 
solution,  to  add  several  cc.  of  concentrated  HC1  ,  heat  to  boiling,  and  pass  in  H2S 
for  some  time  to  precipitate  the  rest  of  the  As  .  In  the  cold,  H3Asp4  is  very  slow- 
ly precipitated  by  H2S  ,  but  strong  HC1  and  heat  accelerate  the  reaction  very  much. 

It  is  essential  that  the  sulphides  be  thoroughly  washed  before  treatment  with 
HC1. 

CuS  is  slightly  soluble  in  (NH4)2SX  and  may  give  a  coloration  when  the  solution 
is  acidified. 

(NH4)2S  ,  which  is  colorless,  gives  no  precipitate  of  S  upon  addition  of  excess  of 
acid;  (NH4)2SX  ,  yellow,  always  gives  more  or  less  S  ,  white  and  difficult  to  filter. 


2S 


2(NH4)2SX  +  4HC1 
(NH4)2S  +  2HC1  = 

Make  a  blank  test  on  the  Fe  wire  used,  to  see  that  its  solution  in  HC1  gives  no 
test  for  Sn  with  molybdate. 

5.  The  sulphides  of  arsenic  are  readily  soluble  in  ammonium  carbonate  (§69, 
5c)  and  are  thus  separated  from  the  sulphides  of  Sb  and  Sn ,  which  are  practically 
insoluble.  The  following  table  suggests  a  method  of  analysis  based  upon  this 
property  of  these  sulphides. 

Digest  the  mixed  sulphides  with  solution  of  ammonium  carbonate  and 
filter. 


Residue:  SnS2  ,  SbaS5  ,  (S)  . 

Dissolve  in  hot  hydrochloric  acid   (5c,   §70 
and  §71). 

Solution:  SnCl4  ,  SbCl8  . 

Treat  with  zinc  and  hydrochloric  acid  in 
Marsh's  apparatus   (§69,  6'a). 


Deposit:   Sn  ,    (Sb)  . 

Dissolve  by  hydro- 
chloric acid. 

Solution:  SnCl2  . 
(Residue,  Sb  .) 

Test  by  ammoniacal 
silver  nitrate  and 
by  mercuric  chlo- 
ride (§71,  6i  and  ;). 


Gas:  SbH3  . 

(Test  the  spots, 
§69,  6'c,  1.) 

Receive  the  gas  in 
solution  of  silver 
nitrate.  Dissolve  the 
precipitate  (Sb  Ags ) 
(§70,  6j),  and  test 
by  H2S  (§87  and 
§89). 


Solution: 
(NH4),AsS4  +  (NH4)3As04 

and 
(NH4)4As2S5  +  (NH4)4As205  . 

Precipitate  by  hydrochloric  acid; 
filter;  wash  the  precipitate  and 
dissolve  it  by  chlorine  gener- 
ated from  a  minute  fragment  of 
potassium  chlorate  and  a  little 
hydrochloric  acid  (§69,  5c). 

Expel  all  free  chlorine  (note  9, 
and  §69,  10). 

Solution:  H8As04  . 

Apply  Marsh's  Test,  as  directed  in 
§69,  fi'rt,  testing  the  spots  (§69. 
6'c) ;  receiving  the  gas  in  solu- 
tion of  silver  nitrate,  and  test- 
ing the  resulting  solution  (§87). 

Examine  the  original  solution,  as 
indicated  in  §88,  1. 


The  arsenic  may  also  be  identified  by  adding  HC1  to  the  ammonium 
carbonate  solution,  passing  H2S  and  dissolving  the  precipitate  in  a  small 
amount  of  concentrated  HN03  .  The  arsenic  will  be  oxidized  to  H3As04  . 
Divide  the  solution  into  two  parts.  Cautiously  neutralize  one  portion  with 


§86,  10.  DIRECTIONS    FOR  ANALYSIS   WITH  NOTES,  121 

ammonia.  When  the  solution  is  nearly  neutral,  add  AgN03  arid  a  drop  or 
two  of  ammonia  without  shaking  the  solution.  A  reddish  brown  ring  of 
Ag3As04  will  form  at  the  neutral  zone  of  the  solution.  To  the  other  por- 
tion, add  magnesia  mixture  and  ammonia  until  the  solution  is  alkaline.  A 
crystaline  precipitate  of  MgNH4As04  will  form  on  standing. 

The  plan  above  given  may  be  varied  by  separating  antimony  and  tin  by  'ammo- 
nium carbonate  in  fully  oxidized  solution,  as  follows:  The  Sb2S5  and  SnS2  are 
dissolved  by  nitrohydrochloric  acid,  to  obtain  the  antimony  as  pyroantimonic 
acid.  The  solution  is  then  treated  with  excess  of  ammonium  carbonate,  in  a 
vessel  wide  enough  to  allow  the  carbonic  acid  to  escape  without  waste  of  the 
solution. 

The  soluble  diammonium  dihydrogen  pyroantimonate,  (NH4)2H2Sb2O7  ,  is 
formed.  Meanwhile  the  SnCl4  is  fully  precipitated  as  H2SnO8  (§71,  6a),  and 
may  be  filtered  out  from  the  solution  of  pyroantimonate. 

The  liability  of  failure,  in  this  mode  of  separating  antimony  and  tin,  lies  in 
the  non-formation  of  pyroantimonic  acid  by  nitrohydrochloric  acid.  The  ordi- 
nary antimonic  acid  forms  a  less  soluble  ammonium  salt,  but  this  acid  is  not 
so  likely  to  occur  in  obtaining  the  solution  with  nitrohydrochloric  as  anti- 
tnonous  chloride,  SbCl3  .  Excess  of  ammonium  carbonate  does  not  redissolve 
the  Sb203  which  it  precipitates  from  SbCls  ,  as  stated  in  §70,  6a. 

The  above  plan  may  also  be  varied  as  follows:  After  removal  of  the  arsenic 
sulphide  with  (NH4)2CO8  ,  the  residue  is  dissolved  in  strong  HC1  ,  not  using 
KC1O3  or  HNO3  .  The  solution  consists  of  SnCl4  and  SbCl,  .  Divide  in  two 
portions:  (1)  Add  Sn  on  platinum  foil.  A  black  precipitate  indicates  Sb°  . 
(2)  Add  iron  wire,  obtaining  Sb°  and  Sn";  filter  and  test  the  filtrate  for  Sn,  by 
Hg-Cl2  (Pieszczek,  Arch.  Pharm.,  1891,  229,  667). 

6.  The  sulphides  of  As  ,  Sb  and  Sn  are  all  decomposed  by  concentrated  nitric 
acid,  which  furnishes  a  basis  of  an  excellent  separation  of  the  arsenic  from  the 
antimony   and  tin    (Vaughan,   American  Chemtet,    1875,    6,    41).     The   sulphides 
reprecipitated  from  the  (NH4)2SX  solution  by  HC1  are  well  washed,  transferred 
to  an  evaporating  dish,  heated  with  concentrated  HN03  until  brown  fumes  are 
no  longer  evolved,  and  then  evaporated  to   dryness,  using  sufficient  heat  to 
expel  the  HN08  and  the  H2S04  formed  by  the  action  of  the  HNO3  upon  the  S  . 
The  heating  should  be  done  on  the  sand  bath.     The  cooled  residue  is  digested 
for  a  few  minutes  with  hot  water,  the  arsenic  passing  into  solution  as  H3AsO4  , 
and  the  antimony  and  tin  remaining  as  residue  of  Sb2Ofl  and  Sn02  .     The  pres- 
ence of  arsenic  may  be  confirmed  by  the  reactions  with  AgN08  (§69,  6;),  CuSO4 
(§69,  Gfc)  by  the  Marsh  test  (§69,  6'a),  or  by  precipitation  with  magnesia  mix- 
ture (§69,  6i).     A  portion  of  the  residue  may  be  tested  in  the  Marsh  apparatus 
for  the  Sb  (§70,  6;),  another  portion  may  be  reduced  and  dissolved  in  an  open 
dish  with  Zn  and  HC1  (not  allowable  if  As  be  present,  §71,  10),  and  the  result- 
ing SnCl2  identified  by  the  reaction  with  HgCl,  (§71,  61). 

7.  The    precipitated    sulphides    must   be    thoroughly    washed    to    insure    the 
removal  of  the  ammonium  salts,  since  in  their  presence  the  dangerously  ex- 
plosive nitrogen  chloride  (§268,  1)  could  be  formed  when  the  sulphides  were 
dissolved  in  HC1  with  the  aid  of  KC1OS  . 

8.  Instead  of  chlorine   (HC1  +  KC1O3),  nitrohydrochloric  acid  may  be  em- 
ployed, but  it  is  liable  to  cause  the  formation  of  a  white  precipitate  of  Sb,05 
and  Sn02  . 

9.  The  chlorine  should  all  be  removed,  as  the  metals  cannot  be  reduced  by 
the  Zn  and  HSSO4  in  the  Marsh  apparatus  in  the  presence  of  powerful  oxidizing 
agents   as  Cl .     This  would  also   require   evaporation   to  expel   the   HNO8  ,   i 
nitrohydrochloric  acid  were  used  to  effect  solution. 

10.  Hydrogen  peroxide,  H2O3  ,  decomposes  the  sulphides  of  arsenic  and  anti- 
mony with  oxidation.     The  arsenic  will  appear  in  the  solution,  the  antimony 
remaining  as  a  white  precipitate  of  the  oxide   (a  sharp  separation)    (Luzzato, 
Arch.  Pharm.,  1886,  224,  772). 


122  DIRECTIONS  FOR   ANALYSIS   WITH  NOTES.  §87. 

§87.  Manipulation. — The  solution  of  the  metals  of  the  tin  group  is 
then  ready  to  be  transferred  to  the  Marsh  apparatus  (the  directions  for 
the  use  of  the  Marsh  apparatus  are  given  under  arsenic  (§69,  6'a),  and 
should  be  carefully  studied  and  observed.  They  will  not  be  repeated 
here).  Only  a  portion  of  the  solution  should  be  used  in  the  Marsh  appar- 
atus, the  remainder  being  reserved  for  other  tests.  The  gas  evolved  from 
the  Marsh  apparatus  is  passed  into  a  solution  of  silver  nitrate,  which  by 
its  oxidizing  action  effects  a  good  separation  between  the  arsenic  and 
antimony  (§89,  2) : 

AsH3  +  6AgN03  +  3H20  =  H3As03  +  6Ag  +  6HNO3 

SbH3  +  3AgN03  =  SbAg3  +  3HNO3 

The  hard  glass  tube  of  the  Marsh  apparatus  is  heated  while  the  gas  is 
being  generated,  a  mirror  of  arsenic  and  antimony  being  deposited,  due 
to  the  decomposition  of  the  gases  (§69,  6'c) :  2SbH3  —  2Sb  +  3H2 .  The 
ignited  gas  is  brought  in  contact  with  a  cold  porcelain  surface  for  the 
production  of  the  arsenic  and  antimony  spots  (§69,  6'&).  Failure  to  obtain 
mirror,  spots,  or  a  black  precipitate  in  the  AgN03  is  proof  of  the  absence 
of  both  arsenic  and  antimony.  The  black  precipitate  obtained  in  the 
silver  nitrate  solution  is  separated  by  filtration,  washed  and  reserved  to  be 
tested  for  antimony.  The  nitrate  is  treated  with  HC1 ,  or  a  metallic 
chloride,  as  CaCl2  or  Nad ,  to  remove  the  excess  of  silver  and,  after  evapor- 
ation to  a  small  volume,  is  precipitated  with  H2S .  A  lemon-yellow  pre- 
cipitate indicates  arsenic.  The  black  precipitate  from  the  silver  nitrate 
solution  is  dissolved  in  hot  reagent  HC1  :  SbAg,  +  6HC1  —  SbCl3  + 
3AgCl .  The  excess  of  acid  is  removed  by  evaporation,  a  little  water  is 
added  (§70,  5d  and  §59,  5c)  and  the  AgCl  removed  by  nitration.  The 
nitrate  is  divided  into  two  portions.  To  one  portion  H2S  is  added;  an 
orange  precipitate  indicates  antimony.  The  H2S  may  give  a  black  precipi- 
tate of  Ag2S  from  the  AgCl  held  in  solution  by  the  HC1 .  If  this  be  the 
case,  to  the  other  portion  one  or  two  drops  of  KI  are  added  and  the 
solution  filtered.  This  filtrate  is  now  tested  for  the  orange  precipitate 
with  H2S . 

The  mirror  obtained  in  the  hard  glass  tube  should  be  examined  as 
directed  in  the  text,  especially  by  oxidation  and  microscopic  examination 
(§69,  6'c  5).  The  spots  should  be  tested  with  NaClO  and  by  the  other  tests 
as  given  in  the  text  (§69,  6'c  1). 

§88.  Notes. — Arsenic. — 1.  All  compounds  of  arsenic  are  reduced  to  arsine  by 
the  Zn  and  H2SO4  in  the  Marsh  apparatus.  Hence  if  strong-  oxidizing  agents 
are  absent,  the  original  solution  or  powder  may  be  used  directly  in  the  Marsh 
apparatus  for  the  detection  of  arsenic;  but  sulphides  should  not  be  present. 

2.  The  burning  arsine  forms  As.,03  ,  which  may  be  collected  as  a  heavy  white 
powder  on  a  piece  of  black  paper  placed  under  the  flame.  Antimony  will  also 
deposit  a.  similar  heavy  white  powder, 


§90.  DIRECTIONS  FOR  ANALYSIS    WITH  NOTES.  123 

3.  The  arsine    evolved  is  not    decomposed   (faint   traces    decomposed)    upon 
passing-  through  a  drying1  tube  containing-  soda  lime  or  through  a  solution  of 
KOH  (distinction  and  separation  from  antimony). 

4.  Arsenites  and  arsenafes  are  distinguished  from  each  other  by  the  following' 
reactions:  (a)  Arsenous  acid  solution  acidulated  with  HC1  is  precipitated  in  the 
cold  instantly  by  H2S;   arsenic   acid  under   similar  conditions   is  precipitated 
exceedingly  slowly  (§69,  6e).     (b)  Neutral  solutions  of  arsenites  give  a  yellow 
precipitate  with  AgN03;  neutral  solutions  of  arsenates  give  a  brick-red  pre- 
cipitate.    Both  precipitates   are  soluble  in   acids   or  in   ammonium  hydroxide 
(§59,  6#).     (c)  Magnesia,  mixture  precipitates  arsenic  acid  as  white  magnesium 
ammonium  arsenate,  MgNH4AsO4;  no  precipitate  with  arsenous  acid  (§189,  60). 

(d)  HI  gives  free  iodine  with  arsenic  acid;  not  with  arsenous  acid   (§69,  6f). 

(e)  Alkaline  solutions  of  arsenous  acid  are  immediately  oxidized  to  the  pentad 
arsenic  compounds  by  iodine  (§69,  10).     (/)  Potassium  permanganate  is  imme- 
diately decolored  by  solutions  of  arsenous  acid  or  arsenites;  no  reaction  with 
arsenates  (§69,  10). 

§89.  Notes. — Antimony. — 1.  If  antimony  be  present  in  considerable  amount, 
it  (in  the  form  of  the  sulphide)  is  most  readily  separated  from  arsenic  by 
boiling-  with  strong-  HC1  (solution  of  the  antimony  sulphide,  (§70,  6e));  or  by 
dig-esting  with  (NH4)2C03  or  NH4OH  (solution  of  the  arsenic  (§69,  5c)). 

2.  For 'the  detection  of  traces  of  antimony,  the  most  certain  test  is  in  its 
volatilization  as  stibine  in  the  Marsh  apparatus  and  precipitation  as  SbAg3  , 
antimony  argentide,  with  AgN03;  this  is  a  good  separation  from  arsenic  and 
tin,  and  after  filtration  it  remains  to  dissolve  the  SbAg3  in  concentrated  HC1 
and  identify  the  Sb  as  the  orange  precipitate  of  Sb2S3  .     The  formation  of  the 
black  precipitate  in  the  AgNO3  solution  must  not  be  taken  as  evidence  of  the 
presence  of  antimony,  as  arsine  gives  a  black  precipitate  of  metallic  silver  with 
AgNO3  .     A  trace  of  antimony  may  be  found  in  the  filtrate  from  the  SbAgs  , 
hence  a  slight  yellow-orange  precipitate  from  this  solution  must  not  be  taken 
as  evidence  of  arsenic  without  further  examination  (§69,  7). 

3.  Sb,S3  is  precipitated  from  solutions  quite  strongly  acid  with  HC1 ,  i.  e.,  in 
the  presence  of  equal  parts  of  the  concentrated  acid  (sp.  gr.  1.20).     Tin  is  not 
precipitated  as  sulphide  if  there  be  present  more  than  one  part  of  the  con- 
centrated acid  to  three  of  the  solution  (§70,  6€).     This  is  a  convenient  method 
0f  separation.     The  addition  of  one  volume  of  concentrated  HC1  to  two  volumes 
of  the  solution  under  examination  before  passing-  in  the  H2S  will  prevent  the 
precipitation  of  the  tin  while  allowing-  the  complete  precipitation  of  the  anti- 
mony. 

4-  If  the  sulphides  of  As  ,  Sb  and  Sn  are  evaporated  to  dryness  with  con- 
centrated HN03;  the  residue  strongly  fused  with  Na2C03  and  NaOH;  and  the 
cooled  mass  disintegrated  with  cold  water,  the  nitrate  will  contain  the  arsenic 
as  sodium  arsenate,  Na3As04  ,  and  the  tin  as  sodium  stannate,  Na2SnO3;  while 
the  antimony  remains  as  a  residue  of  sodium  pyroantimonate.  Na2H2Sbo07 
(§70,  7). 

5.  Stibine  is  evolved  much  more  slowly  than  arsine  in  the  Marsh  apparatus, 
and  some  metallic  antimony  will  nearly  always  be  found  in  the  flask  with  the 
tin  (§70,  6;). 

6.  If  organic  acids,  as  tartaric  or  citric,  be  present,  they  should  be  removed 
by  careful  ig-nition  with  K2CO3  as  preliminary  to  the  preparation  of  the  sub- 
stance for  analysis,  since  they  hinder  the  complete  precipitation  of  the  anti- 
mony with  H2S  (§70,  6e). 

7.  Antimonic  compounds  are  reduced  to  the  antimonous  condition  by  HI  with 
liberation  of  iodine   (§70,  Qf  and  10).     Chromates  oxidize  antimonous  salts  to 
antimonic  salts  with  formation  of  green  chromic  salts  (§70,  6ft).     KMnO4  also 
oxidizes  antimonous  salts  to  antimonic  salts,  a  mang-anous  salt  being  formed 
in  acid   solution    (§70,   6ft).     No   reaction   with   antimonic   salts.     Antimonous 
salts  reduce  gold  chloride;  antimonic  salts  do  not  (§73,  10). 

§90.  Manipulation. — The  contents  of  the  generator  of  the  Marsh  appar- 
atus should  be  filtered  and  washed.     The  nitrate,  if  colorless,  may  be 


124  DIRECTIONS  FOB  ANALYSIS    WITH  NOTES.  §91,^. 

rejected  (absence  of  Mo).  A  colored  filtrate,  blue  to  green-brown  or  black, 
indicates  the  probable  presence  of  some  of  the  lower  forms  of  molybdenum. 
The  solution  should  be  evaporated  to  dryness  with  an  excess  of  HN03  , 
which  oxidizes  the  molybdenum  to  molybdic  acid,  Mo03 .  The  residue 
is  dissolved  in  NH4OH  (the  zinc  salt  present  does  not  interfere)  and  poured 
into  moderately  concentrated  nitric  or  hydrochloric  acid  (§75,  6d  footnote). 
This  solution  is  tested  for  molybdenum  by  Na2HP04 .  The  original  solu- 
tion should  also  be  examined  for  the  presence  of  molybdenum  as  molybdic 
acid  or  molybdate  (§75,  6d). 

The  residue  from  the  generator  of  the  Marsh  apparatus  may  contain 
Sb ,  Sn ,  Ail ,  and  Pt  with  an  excess  of  Zn .     It  should  be  dissolved  as 
much  as  possible  in  HC1 .     Sb ,  Au ,  and  Pt  are  insoluble  (§70,  5a).     The 
Sn  passes  into  solution  as  SnCl2  and  gives  a  gray  or  white  precipitate  with 
HgCl2 ,  depending  on  amount  of  the  latter  present  (§71,  6;) : 
SnCl2  +  HgCla  —  SnCl4  -f  Hg 
SnCl,  +  2HgCl2  =  2HgCl  +  SnCl4 

The  presence  of  Sn"  should  always  be  confirmed  by  its  action  in  fixed 
alkali  solution  upon  an  ammoniacal  solution  of  AgN03 ,  giving  Ag° 
(§71,  6i). 

Au  and  Pt  may  be  detected  in  the  residue,  but  it  is  preferable  to  precipi- 
tate them  from  a  portion  of  the  original  solution  by  boiling  with  ferrous 
sulphate  (6/&,  §§73  and  74).  Both  metals  are  precipitated.  They  are  then 
dissolved  in  nitro-hydrochloric  acid  and  evaporated  to  dryness  with  am- 
monium chloride  on  the  water  bath.  The  residue  is  treated  with  alcohol 
which  dissolves  the  double  chloride  of  gold  and  ammonium,  leaving  the 
platinum  double  salt  as  a  precipitate,  which  is  changed  to  the  metal  upon 
ignition.  The  alcoholic  solution  is  evaporated,  taken  up  with  water  and 
the  gold  precipitated  by  treating  with  FeS04  (§73,  6h),  by  boiling  with 
oxalic  acid  (§73,  6&),  or  by  treating  with  a  mixture  of  SnCl2  and  SnCl4 
(Cassius'  purple)  (§73,  6#). 

If  a  portion  of  the  original  solution,  free  from  HN03 ,  be  boiled  with 
oxalic  acid  the  gold  is  completely  precipitated  as  the  metal,  separation 
from  the  platinum  which  is  not  precipitated  (§74,  6&). 

§91.  — Notes. — Molybdenum. — 1.  In  the  regular  course  of  analysis,  molyb- 
denum remains  in  the  flask  of  the  Marsh  apparatus  as  a  dark  colored  solution, 
the  Zn  and  H2SO4  acting-  as  a  reducing  agent  upon  the  molybdic  acid. 

2.  If  the  molybdenum  be  present  in  solution  as  molybdic  acid  or  a  molybdate, 
it  may  be  separated  in  the  acid  solution  from  the  other  metals  by  phosphoric 
acid  in  presence  of  ammonium  salts,  forming  the  ammonium  phosphomolyb- 
date;  insoluble  in  acids,  but  soluble  in  ammonium  hydroxide  (§75,  6d). 

3.  In  ammoniacal  solution  of  a  phosphomolybdate,  magnesium  salts  precipi- 
tate the  phosphoric  acid,  leaving  the  molybdenum  as  ammonium  molybdate  in 
solution,  which  may  be  evaporated  to  cr3rstallization    (method   of  recovering 
ammonium  molybdate  from  the  ammonium  phosphomolybdate  residues). 

§92.  Tin. — 1.  Tin  requires  the  presence  of  much  less  HC1  to  prevent  its  pre- 
cipitation by  HoS  than  arsenic  or  antimony  (§89,  3). 


§96.  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES.  125 

2.  The  yellow  ammonium  sulphide  (NH4).,SX  must  be  used  to  effect  solution 
if   tin    (Sn")    be   present,   SnS   being-  practically  insoluble  in   the   normal   am- 
monium sulphide  (§71,  5c). 

3.  Tin  in  the  stannous  condition,  dissolved  in  the  fixed  alkalis    (stannites), 
readily  precipitates   metallic   silver   black   from   solutions   of   silver  salts.     An 
arsenite   (hot)   or  an  antimonite  in  solution  of  the  fixed  alkalis  produces  the 
same  result,  but  not  if  the  silver  salt  be  dissolved  in  a  great  excess  of  ammo- 
nium hydroxide    (§7O,  6i).     This   reaction   also   detects   stannous   salts   in   the 
presence  of  stannic  salts. 

4-  Tin  in  the  Marsh  apparatus  is  reduced  to  the  metal,  and  then  by  solution 
of  the  residue  in  HC1 ,  forms  SnCL  ,  which  may  be  detected  by  the  reduction 
of  HgCl2  to  Hg-Cl  or  Hg°  (§71,  6;),  and  by  the  action  in  fixed  alkali  solution 
upon  the  strong-  ammoniacal  solution  of  silver  oxide  (§71,  6i). 

5.  If  the  Zn  in  the  Marsh  apparatus  is  completely  dissolved,  the  Sn  must  be 
looked  for  in  the  solution,  which  in  this  case  must  not  be  rejected.     The  tin 
remains  as  the  metal  as  long-  as  zinc  is  present  (§135,  10). 

6.  The  presence   of  the  tin  may  be  confirmed  by  its  action  as  a  powerful 
reducing  agent  (§71,  10).     If  it  be  present  as  Sniv  ,  these  tests  must  be  made 
after  reduction  in  the  Marsh  apparatus  or  in  an  open  dish  with  zinc  and  HC1 . 

§93.  Gold. — 1.  Gold"  will  usually  be  met  with  in  combination  with  other  metals 
as  alloys,  and  is  separated  from  most  other  metals  by  its  insolubility  in  all 
acids  except  nitrohydrochloric  acid. 

2.  If  more  than  25  per  cent  of  gold  be  present  in  an  alloy,  as  with  silver, 
the  other  metal  is  not  removed  by  nitric  acid   (§73,  5a).     Either  nitrohydro- 
chloric acid  must  be  used  or  the  alloy  fused  with  about  ten  times  its  weight  of 
silver  or  lead,  and  this  alloy  dissolved  in  nitric  acid  when  the  gold  remains 
behind. 

3.  If  the  presence  of  gold  is  suspected  in  the  solution,  it  should  be  precipi- 
tated with  FeS04  before  proceeding  with  the  usual  method  of  analysis. 

4.  If  gold  be  present  (in  the  usual  method  of  analysis)   it  will  remain  as  a 
metallic  residue  in  the  Marsh  apparatus,  insoluble  in  HC1  and  may  be  identi- 
fied by  the  reactions  for  Au°  . 

5.  The  reactions  of  gold  chloride  with  the  chlorides  of  tin  forming  Cassius' 
purple  (§73,  G<?)  is  one  of  the  most  characteristic  tests  for  gold. 

§94.  Platinum. — 1.  Notes  1  to  4  under  gold  apply  equally  well  for  platinum, 
except  that  it  is  necessary  to  boil  with  FeSO4  to  insure  complete  precipitation 
of  the  platinum. 

2.  Oxalic  acid  is  the  best  reagent  for  the  separation  of  gold  from  platinum 
(§73,  66). 

3.  The  most  important  problems  in  the  analysis  of  platinum  consist  in  its 
separation  from  the  other  metals  of  the  platinum  ores  (§74,  3). 

§96.  Manipulation. — The  well  washed  residue  after  digesting  the  pre- 
cipitated sulphides  of  the  second  group  (the  Tin  and  Copper  Group)  in 
(NH4)2SX  may  contain  any  of  the  metals  of  the  Copper  Group,  and  in 
addition  frequently  contains  sulphur.,  formed  by  the  action  of  the  H2S 
upon  oxidizing  agents :  4FeCl3  +  2H2S  =  4FeCl2  +  4HC1  +  S2  .  Pierce 
the  point  of  the  filter  with  a  small  stirring  rod  and,  with  as  little  water  as 
possible,  wash  the  precipitate  into  a  test-tube,  beaker,  or  small  casserole, 
Sufficient  reagent  nitric  acid  (§324)  should  be  added  to  make  about 
one  part  of  the  acid  to  two  parts  of  water  and  the  mixture  boiled  vigor- 
ously for  two  or  three  minutes :  * 

2Bi2Ss  +  16HN03  :  :  4Bi(NO»).  -f  4NO  +  8H20  -f  38, 
6CdS  -f  1<3HN08  =3  6Cd(N03)2  +  4NO  -f  Oii20  +  38, 

*  If  preferred  the  precipitate  on  the  filter  may  be  washed  with  the  boiling  hot  nitric  acid  of 
the  above  mentioned  strength,  pouring:  the  same  acid  back  upon  the  precipitate,  reheating-  eacJ> 
time,  until  no  further  acflon  takes  place, 


126 


TABLE  FOR  ANALYSIS   OF  THE  COPPER   GROUP. 


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DIRECTIONS   FOR  ANALYSIS   WITH  NOTES.  §97,  1. 

Mercuric  sulphide  is  unattacked  (§58,  Qe)  and  remains  as  a  black  pre- 
cipitate together  with  some  sulphur  as  a  yellow  to  brown-black  precipitate. 
The  precipitate  is  filtered  and  washed  with  a  small  amount  of  hot  water. 
The  filtrate  is  set  aside  to  bo  tested  later,  and  the  black  residue  on  the 
filter  is  dissolved  in  nitro-hydrochloric  acid :  2HgS  +  2C12  =  2HgCl2  +  32 . 
This  solution  is  boiled  to  expel  all  chlorine  and  the  presence  of  mercury 
determined  by  reduction  to  HgCl  or  Hg°  by  means  of  SnCl,  (§58,  60) : 
HgCL  +  SnCl2  ==  Hg  +  SnCl4P  2HgCl2  +  SnCl2  ==  2HgCl  +  SnCl,';  or 
by  the  deposition  of  a  mercury  film  on  a  strip  of  bright  copper  wire 
(§5C,  10) :  HgCl2  +  Cu  =  Hg  +  CuCl2 .  Confirm  further  by  bringing  in 
contact  with  iodine  in  a  covered  dish:  Hg  +  I2  =  HgI2  (Jannaesch,  Z. 
anorg.,  1896,  12,  143).  The  mercury  may  also  be  detected  by  using 
NH4OH  and  KI  as  the  reverse  of  the  Nessler's  test  (§207,  6fc)  (delicate 
1  tj>  31,000)  (Klein,  Arch.  PJiarm.,  1889,  227,  73). 

§97.  Notes. — 1.  The  concentration  of  HNO3  (1-2)  is  necessary  for  the  solution 
of  the  sulphides  of  Pb  ,  Bi  ,  Cu  and  Cd  ,  and  may  also  dissolve  traces  of  HgS  . 
However,  the  concentrated  HNO:i  (sp.  gr.,  1.42)  dissolves  scarcely  more  than 
traces  of  HgS  (§58,  6e).  Long-continued  boiling  of  HgS  with  concentrated 
HNO3  changes  a  portion  of  the  HgS  to  Hg(N03)t.HgS  ,  a  white  precipitate, 
insoluble  in  HNO3  . 

2.  In  the  use  of  nitrohydrochloric  acid  to  dissolve  the  HgS  ,  the  HC1  should 
be  used  in  excess  to  insure  the  decomposition  of  the  nitric  acid,  which  would 
interfere  with  the  reduction  tests  with  SnCl,   and  Cu°  .     One  part  of  HNO3 
to  three  parts  HC1  gives  about  sufficient  HC1  to  decompose   all   the  HN03  , 
hence  in  this  reaction  a  little  more  than  that  proportion  of  HC1  should  be 
u«ed. 

3.  A  small  amount  of  black   residue   left   after   boiling  the   sulphides  with 
HNO3    may   consist   entirely   of    sulphur,    which    can   best    be    determined    by 
burning   the   residue   on    a   platinum   foil   and    noting   the    appearance   of   the 
flame,  the  odor,  and  the  disappearance  of  the  residue.     The  residue  of  sulphur 
frequently  possesses  the  property  of  elasticity  (§256,  1). 

4-  Boiling  the  sulphides  of  the  copper  group  with  HNOS  will  always  oxidize 
a  trace  at  least  of  sulphur  to  H2S04  (§256,  6B,  2),  which  will  form  PbSO4  if  any 
lead  be  present: 

S2  +  4HNC-3  =  2H2S04  +  4NO 
3PbS  +  8HNO3  =  3PbS04  +  4H20  -f-  8NO 

If  the  boiling  ,be  not  continued  too  persistently,  the  amount  of  PbS04  formed 
is  soluble  in  the  HN03  present  (§57,  5o),  and  does  not  at  all  remain  behind 
with  the  HgS  . 

6.  Even  if  only  1  to  2  mg.  of  As  or  Sb  are  present  with  a  large  quantity  (500  ing.) 
of  an  element  of  the  copper  group  enough  is  dissolved  by  either  (NH4)2S  or 
(NH4)2SX  for  the  detection  of  these  metals.  If  only  3  to  5  mg.  of  Sn  are  present 
with  a  large  quantity  of  elements  of  the  copper  group,  all  of  the  tin  may  remain 
undissolved.  When  Cd  is  present  and  the  tin  is  in  the  stannous  state  as  much 
as  15  mg.  of  Sn  may  remain  undissolved  even  in  the  polysulphide  (A.  A.  Noyes, 
J.  Am.  Chem.  Soc.  29,  170).  The  Sn  or  Sb  (present  on  account  of  an  insuffi- 
ciency of  (NH4)2SX)  will  appear  as  a  white  precipitate  mixed  with  the  black 
precipitate  of  HgS  ,  due  to  the  fact  that  HNO3  decomposes  the  sulphides  of  Sb 
and  Sn  ,  forming  the  insoluble  Sb>O5  and  SnO.  : 

6Sb2S3  -f-  20HNO3  =  6Sb2O6  +  9S2  -f  20NO  +  10H2O 

If  these  metals  have  not  been  detected  this  precipitate  must  be  tested.  Aftei 
testing  for  mercury  in  a  portion  of  the  precipitate,  the  paper  may  be  burned 


§100.  DIRECTIONS   FOR  ANALYSIS   WITH  NOTES.  129 

in  a  porcelain  crucible  and  the  residue  fused  with  sulphur  and  sodium  carbonate 
in  the  covered  crucible.  The  tin  and  antimony  will  be  converted  into  soluble 
thio  salts  and  tested  for  according  to  §  84. 

6.  Traces  of  mercury   may   be  detected   by  using-  a   tin-gold   voltaic  couple. 
The  Hg  deposits  on  the  Au  .  and  can  be  sublimed  and  identified  with  iodine 
vapor.     Arsenic  gives  similar  results  (Lefort,  C.  r.,  1880,  90,  141). 

7.  Merciiry  may  quickly  be  detected  from  all  of  its  compounds  by  ignition 
in  a  hard  glass  tube  with  fusion  mixture  (Na,COs  +  K,C03)   (§58,  7),  and  then 
adding-  a  few  drops  of  HN03  (concentrated)  and  a  small  crystal  of  KI  .     Upon 
warming-  the  iodine  sublimes  and  combines  with  the  sublimate  of  Hg  ,  forming 
the  scarlet  red  HgI2  .     As  and  Sb  both  give  colored  compounds  with  iodine,  de- 
composed by  HNO3'(Johnstone,  C.  N.,  1889,  59,  221). 

§98.  Manipulation. — To  the  filtrate  containing  the  nitric  acid  solution 
of  the  sulphides  of  Pb ,  Bi ,  Cu ,  and  Cd ,  should  be  added  about  two  cc.  of 
concentrated  H2S04  and  the  mixture  evaporated  on  a  sand  bath  or  over 
the  naked  flame  in  a  casserole  or  evaporating  dish  until  the  fumes  of 
H2S04  are  given  off: 

Pb(NOs)2  +  H,,S04  =  PbS04  +  2HNO3 

Cu(N03),  +  H2S04  =  CuS04  +  2HN03 

About  20  cc.  of  50  per  cent  alcohol  should  be  added  to  the  well  cooled 
mixture  and  the  whole  transferred  to  a  small  glass  beaker.  Upon  giving 
the  beaker  a  rotatory  motion  the  heavy  precipitate  of  PbS04  will  collect 
in  the  center  of  the  beaker,,  and  its  presence  even  in  very  small  amounts 
may  be  observed.  The  filtrate  from  the  PbS04  should  be  decanted  through 
a  wet  filter,  and  the  PbS04  in  the  beaker  may  be  further  identified  by  its 
transference  into  the  yellow  chromate  with  K2Cr04  or  into  the  yellow  iodide 
with  KI  (57,  G/  and  h). 

§99.  Notes. — 1.  In  analysis,  if  lead  was  absent  in  the  silver  group,  it  is 
advantageous  to  test  only  a  portion  of  the  nitric  acid  solution  with  H^SO4  for 
lead,  and  if  that  metal  be  not  present,  the  above  step  may  be  omitted  with 
the  remainder  of  the  solution  and  the  student  may  proceed  at  once  to  look 
for  Bi  ,  Cu  and  Cd  .  If,  however,  lead  is  present,  the  whole  of  the  solution 
must  be  treated  with  H,S04  . 

2.  The  nitric  acid  should  be  removed  by  the  evaporation,  as  PbSO4  is  quite 
appreciably  soluble  in  HNO3  (§57,  5c). 

3.  The  H,SO4  should  be  present  in  some  excess,  as  PbSO4  is  less  soluble  in 
dilute  H;,S04  than  in  pure  water  (§57,  5c). 

4.  Alcohol  should  be  present,  as  it  greatly  decreases  the  solubility  of  PbS04 
in  water  or  in  dilute  H2SO4  (§57,  5c,  Ge). 

5.  Too  much  alcohol  must  not  be  added,  as  sulphates  of  the  other  metals 
present  are  also  less  soluble  in  alcohol  than  in  water   (§77,  5c).     These  sul- 
phates, if  precipitated  by  the  alcohol,  are  readily  dissolved  on  dilution  with 
water. 

6.  If  the  (NH4)2SX  had  not  been  well  removed  by  wrashing,  ammonium  sul- 
phate  would    be   present   at   this    point,    greatly    increasing   the    solubility   of 
PbS04  (§57,  5c). 

§100.  Manipulation.— The  filtrate  from  the  PbS04  should  be  boiled  to 
expel  the  alcohol  (or  if  Pb  be  absent  evaporate  the  nitric  acid  solution  of 
division  B)  and  then  carefully  neutralized  with  NH4OH .  An  excess  of 


130  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES.  §101, 1. 

NH4OH  should  be  added  to  dissolve  the  precipitates  of  Cu(OH)2  and 
Cd(OH)2 ,  leaving  the  Bi(OH)3  as  a  white  precipitate.  The  solution  should 
be  filtered,  the  precipitate  thoroughly  washed,  and  then  treated  upon  the 
filter  with  a  hot  solution  of  potassium  stannite,  K2Sn02 .  A  black  pre- 
cipitate is  evidence  of  the  presence  of  Bi  (§76,  Qg). 

§101.  Notes. — 1.  If  the  precipitate  of  the  sulphides  of  the  second  group  was 
not  well  washed,  the  hydroxides  of  the  metals  of  the  iron  group  (Al ,  Cr  and 
Fe)  may  be  present  at  this  point.  The  precipitate  of  A1(OH)3  would  be  white, 
but  would  not  give  a  black  precipitate  with  K2SnO,  . 

2.  If  an  insufficient   quantity  of    (NH4)2SX  was   used,   Sb   and  Sn  would   be 
present  and  give  a  white  precipitate  with  the  NH4OH  . 

3.  If  the  lead  had  not  been  removed  it  would  appear  as  a  white  precipitate 
with  the  NH4OH  ,  and  would  give  a  brownish-black  precipitate  with  the  hot 
K2SnO2  (§57,  6#).     The  presence  of  a  permanent  white  precipitate  with  NH4OH 
must  never  be  taken  as  final  evidence  of  the  presence  of  Bi  . 

4.  As  a   confirmatory  test  for  the  presence  of  Bi  ,   a  portion   of  the  white 
precipitate  with  NH4OH  should  be  dissolved  in  HC1  and  the  solution  evapo- 
rated  nearly   to   dryness   to    remove    the    excess    of    HC1 .     Now   upon    adding 
water,  a  white  precipitate  of  BiOCl ,  bismuth  oxychloride,  will  be  obtained  if 
Bi  is  present  (§76,  5(7). 

§102.  Manipulation.— If  the  ammoniacal  filtrate  from  the  Bi(OH)3  is  of 
a  blue  color,  that  is  sufficient  evidence  of  the  presence  of  Cu  unless  nickel 
was  precipitated  in  the  second  group.  In  absence  of  a  blue  color  a  portion 
of  the  solution  should  be  acidulated  with  acetic  acid  and  then  to  this  solu- 
tion a  few  drops  of  potassium  ferrocyanide,  K4Fe(CN)c ,  should  be  added. 
A  brick-red  precipitate  is  evidence  of  copper.  Or  to  the  acidulated  solu- 
tion a  bright  nail  or  piece  of  iron  wire  may  be  added,  obtaining  a  film  of 
metallic  copper.  If  sufficient  copper  be  present  to  give  a  blue  color  to 
the  solution,  before  testing  for  cadmium  a  solution  of  KCN  should  be 
added  until  the  blue  color  disappears.  Then  the  addition  of  H2S  will 
give  a  yellow  precipitate  for  cadmium. 

§103.  Notes. — 1.  The  precipitate  of  the  brick-red  Cu2Fe(CN)6  is  a  much  more 
delicate  test  for  copper  than  the  blue  color  to  the  ammoniacal  solution  (§77, 
66).  Cd  gives  a  white  precipitate,  insoluble  in  the  acid. 

2.  The  student  should   not  forget   that  in  the  regular  course  of  analysis   a 
trace  of  copper  may  be  lost  by  the  solubility  of  the  sulphide  in  (NH4),SX  .     If 
mercury  has  been  shown  to  be  absent,  the  sulphides  of  the  tin  group  (second 
group,  division  A)  should  be  dissolved  by  the  addition  of  a  fixed  alkali  sulphide 
(§71,   6e),  K2S  or  Na,S ,  which  does  not  dissolve   CuS.     In   case  mercury   be 
present,  the  presence  or  absence  of  small  amounts  of  copper  must  be  deter- 
mined by  the  usual   reactions  for  copper  upon   the   original  solution,   having 
due   regard   for   the   possible   interference   of  metals   which   the    analysis   has 
shown  to  be  present. 

3.  Potassium  cyanide,   KCN  ,  in  excess  changes  cupric   salts  to   the   soluble 
double  salt  of  cuprous  cyanide  and  potassium  cyanide,  K3Cu(CN)4  ,  which  is 
colorless  and  not  precipitated  by  sulphides.     With  cadmium  salts  the  soluble 
double    cyanide.    K,Cd(CN)4  ,    is    formed,    which    is    decomposed    by    sulphides 
forming  CdS  ,  yellow. 

4-  If  preferred,  the  sulphides  of  Cu  and  Cd  may  be  precipitated  from  the 
ammoniacal  solution  by  H2S  and  then  the  black  CuS  dissolved  with  KCN  , 
leaving  a  yellow  precipitate  of  CdS  . 


§104,  6.  RUTHENIUM.  131 

5.  Copper  and  cadmium   may  be  separated  from  each  other  by  reduction  of 
the  copper    (from  the  ammoniacal  solution  acidulated  with  HC1)   with  SnCL 
(§77,  10):  2CuCL  +  SnCl,  =  2CuCl  -f  SnCl4  ,  and  its  precipitation  with  milk 
of  sulphur  (§77,  Ge),  forming-  Cu.S  ,  removal  of  the  tin  with  NH4OH  and  the 
precipitation  of  the  cadmium  with  H2S  . 

6.  From  Ihe  solutions  of  copper  and   cadmium  acidulated  with  HC1  ,  a  hot 
solution  of  Na2S2O3  precipitates  the  copper  as  Cu,S  (§77, Ge),  while  the  cadmium 
remains  in  solution.     From  this  solution  the  cadmium  is  detected  as  the  sul- 
phide by  neutralization  with  NH.OH  and  precipitation  with  H,S  or  (NH4),S. 

7.  The  ammoniacal  solution  of  Cu  and  Cd  may  be  precipitated  with  H,S  , 
and  the  resulting-  sulphides,  after  filtering  and  washing-,  boiled  with  hot  dilute 
H2SO4    (one  of  acid  to  five  of  water).     In  this  solution  the  CuS   (§77,   5c)   is 
unattacked  while  the  CdS  is  dissolved.     The  filtrate  upon  dilution  with  water 
gives  the  yellow  CdS  with  H2S  or  (NH4)2S  (§78,  6e). 


BARER  METALS  OF  THE  TIN  AND  COPPER  GROUP. 
(Second  Group.) 

Ruthenium,  Rhodium,  Palladium,  Iridium,   Osmium,  Tungsten,  German- 
ium, Tellurium,  Selenium. 

§104.     Ruthenium.     Ru  =  101.7.     Valence  two  to  eight. 

1.  Properties.— Specific  gravity,  11.0  to  11.4  (Deville  andDebray,  C.r.,  1876,  83, 
926).     Melting-point  2450°  ?  (Cir.  B.  S.t  35, 1915).     Next  to  osmium  it  is  the  most 
difficultly  fusible  of  all  the  platinum  metals.     A  black  powder  or  a  grayish-white 
crystalline  brittle  metal. 

2.  Occurrence. — In  small  quantities  in  platinum  ores. 

3.  Preparation. — Ignite  the  Pt  residues  in  a  stream  of  chlorine  in  presence  of 
NaCl .     Dissolve  the  fused  mass  in  H2O  ,  add  KN02  ,  neutralize  with  Na2C03  , 
evaporate   to   dryness   and  extract   the  double   nitrites   with   absolute   alcohol 
(separation  from  rhodium).     Add  water  to  the  solution,  distill    off  the  alcohol, 
add  HC1  and  obtain  a  red  solution  of  potassium  ruthenium  chloride.     This  is 
changed   to  the   double   ammonium   salt   and   then   precipitated   with    HgCl2  , 
which  upon  recrystallization  and  ignition  gives  pure  Ru  (Gibbs,  Am.  8.,  1862, 
(2),  34,  349  and  355). 

4.  Oxides    and    Hydroxides. — The    hydroxides,    Ru(OH)2  ,    Ru(OH)3  ,    and 
Ru(OH)4  ,  are  precipitated  from  the  respective  chlorides  by  KOH  .     They  are 
dark   brown    to   black.     Perruthenic    anhydride    or    acid,    RuO4  ,    is    a   golden 
yellow  crystalline  powder,  volatile  even  at  ordinary  temperatures.     It  has  a 
peculiar  odor,  somewhat  like  ozone,  is  sparingly  soluble  in  water,  melts  at  50° 
and  boils  at  a  little  over  100°  (Deville  and  Debray,  B.,  1875,  8,  339).     It  is  pre- 
pared  by   heating  K2RuCl5   with   KOH   into  which   a   current   of  chlorine   is 
passed  or  by  distillation  of  a  Ru  salt  with  KC103   and  HC1 .     The  vapor  is 
yellow  and  is  strongly  irritating  to  the  membrane  of  the  throat. 

5.  Solubilities. — Ru  is  soluble  with  difficulty  in  nitrohydrochloric  acid,  in- 
soluble by  fusion  witk  KHSO4  ,  but  is  soluble  by  fusion  with  KOH  ,  especially 
in    presence    of    KNO3  .     Soluble   in    chlorine,    forming'    a    mixture    of    RuCl2  , 
RuCl3  ,  and  RuCl4  .     The  double  nitrites  are  soluble  in  water  and  alcohol  (sepa- 
ration from  rhodium). 

6.  Reactions. — The  alkalis  precipitate  from  ruthenic  chloride  the  dark  yellow 
hydroxide,  soluble  in  acids,  insoluble  in  the  fixed  alkalis,  soluble  in  NH4OH 
with    a    greenish-brown    color.     H2S    precipitates    slowly    the    black    sulphide 
(formed  at  onc«  by   (NH4)2S),  the  solution  becoming  blue.     The  sulphide  is 
insoluble  in  alkali  ^sulphides.     KI  gives  with  hot  solutions  a  black  precipitate 


132  RHODIUM.  §104,  7. 

of  ruthenic  iodide.  KCNS  forms,  after  some  time  in  the  cold,  a  red  coloration, 
which  upon  heating*  assumes  a  beautiful  violet  color  (characteristic).  The 
double  nitrites  are  soluble,  and  if  to  the  solution  (NH^nS  be  added,  a  char- 
actetristic  crimson  red  liquid  is  obtained.  Upon  standing  the  solution  becomes 
brown,  or  a  brown  precipitate  is  caused  by  excess  of  the  (NH4)2S  . 

7.  Ignition. — If  Ru04  be  heated  to  a  dull-red  heat  the  violet-blue  dioxide  is 
formed  (Debray  and  Joly,  C.  r.,  1888,  106,  328). 

8.  Detection. — By  oxidation  and  distillation  as  Ru04  . 

9.  Estimation. — Keduced  to  the  metal  and  weighed  as  such. 

10.  Oxidation.— Ru04    heated    with    HC1    forms    RuCl3  ,    evolving    chlorine. 
B,U8  —  x  is  changed  to  Bu04  by  distilling  with  KC103  and  HC1 .     Zn  reduces 

Ru  solutions  to  the  metal,  with  an  indigo-blue  color  during  transition  from 
Ruiv  to  Run  . 


§105.  Rhodium.     Rh  =  102.9  .     Valence  two,  three  and  four. 

1.  Properties.— Specific  grarity,  12.1  (Deville  and  Debray,  C.  r.,  1874,  78,  1782). 
Melting  point,  1950°  (Cir.  B.  S.,  35,  1915).     It  is  a  white  metal,  nearly  as  ductile 
and  malleable  as  Ag.     The  metal  precipitated  by  alcohol  or  formic  acid  appears 
as  a  black  spongy  mass  (Wilm,  B.,  1881,  14,  629). 

2.  Occurrence. — Found  in  platinum  ores. 

3.  Preparation. — Fusion  of  the  Pt  residues  with  Pb  ,  digestion  with  HN03 
and  then  Cl  ,  converting  the  Rh  into  the  chloride,  from  which  solution  it  is 
precipitated    as   the   double    ammonium    chloride    by    fractional    precipitation. 
See  Gibbs  (/.  pr.,  1865,  94,  10)  and  Wilm  (5.,  1883,  16,  3033). 

4.  Oxides   and   Hydroxides. — Rh(OH)3    is    precipitated    from    a    solution    of 
sodium  rhodium  chloride  by  an  excess  of  KOH  .     It  is  a  black  gelatinous  pre- 
cipitate, forming  the  oxide  upon  ignition.     Rhodium  fused  with  KOH  and  KNO3 
gives  Rh02  ,  a  brown  powder,  insoluble  in  acids  or  alkalis. 

5.  Solubilities. — The  pure  metal  or  the  alloy  with  Au  or  Ag  is  almost  in- 
soluble in  acids;  alloyed  with  Bi  ,  Pb  ,  Cu  or  Pt ,  it  is  soluble  in  HNO3  (Deville 
and  Debray,  I.e.).     Attacked  by  chlorine  the  most  easily  of  all  the  Pt  metals. 
The  precipitated  metal,  a  gray  powder,  is  soluble  in  HC1  in  presence  of  air  to 
a  cherry-red  color. 

6.  Reactions.— Alkali  hydroxides  and  carbonates  precipitate  solutions  of  Rh 
salts   as   Rh(OH)3  ,   yellow,   insoluble   in   acids,   soluble   in   excess   of   NH4OH  , 
forming  a  rhodium  ammonium  base,  precipitated  by  HC1  as  a  bright  yellow 
crystalline     salt,     chloro-purpureo-rhodium     chloride,     Rh(NH3)5Cl3  .      Alkali 
nitrites  precipitate  alcoholic  solutions  of  rhodium  chloride  as  alkali-rhodium 
nitrite  (Gibbs,  Am.  S.,  1862,  (2),  34,  341)    (separation  from  ruthenium).     From 
a  hot  solution  of  Rh  salt,  H2S  precipitates  the  sulphide,  insoluble  in  the  alkali 
sulphides:  the  sulphide  precipitated  from  the  cold  solution  is  soluble  in  alkali 
sulphides.     KI  precipitates  from  hot  solutions  a  black-brown  rhodium  iodide. 

7.  Ignition. — When  the  metal  or  its  compounds  are  repeatedly  fused  with 
HPO3    or   KHS04  ,  the  corresponding  Rh  salts   are  formed.     The  mass  fused 
with  KHS04  is  soluble  in  water  to  a  yellow  color,  turning  red  with  HC1  . 

8.  Detection. — By  ignition  as  given  above.     Also  to  the  concentrated  neutral 
solution   add   fresh    NaCIO   solution.     To   the   yellow   precipitate   add    a    small 
amount   of   HC2H302    and    shake    till    an    orange-yellow    solution    is    obtained. 
After  a   short  time   the   solution   becomes   colorless,   then   a   gray   precipitate 
separates  out  and  the  solution  assumes  a  sky-blue  color  (Demarcay,  C.  r.,  1885, 
101,  951). 

9.  Estimation. — It  is  reduced  to  the  metal  and  weighed  as  such. 

10.  Oxidation.— Solutions  of  rhodium  salts  are  reduced  to  the  metal  by  Zn  . 
All    Rh   compounds   are    reduced   to   the   metal   by   heating   in    a   current   of 
hydrogen. 


§106,  6.  PALLADIUM.  133 

10.  Oxidation. — Solutions  of  rhodium  salts  are  reduced  to  the  metal  by  Zn.  All 
Rh  compounds  are  reduced  to  the  metal  by  heating  in  a  current  of  hydrogen. 

§106.  Palladium.    Pd  =  100.7.     Valence  two  and  four. 

1.  Properties. — Specific  gravity,  11.4  (Deville  and  Debray,  C.  r.,  1857,  44,  1101)' 
Melting  point  1549°  (Cir.  B.  S.,  36,  1915).     It  conducts  electricity  about  one-eighth 
as  well  as  silver  (Matthiessen,  Pogg.,  1858,  103,  428).     Palladium  has  about  the 
color  and  lustre  of  silver.    The  metal  when  only  slightly  heated  assumes  a  rainbow 
tint  from  green  to  violet.    Because  of  its  general  properties,  it  is  to  be  classed  with 
the  platinum  metals,  yet  in  its  reaction  with  acids  it  is  markedly  different.     In  the 
air  at  ordinary  temperature  it  is  but  slightly  tarnished,  but  at  a  red  heat  it  becomes 
covered  with  a  coating  of  the  oxide.    The  finely  divided  metal,  palladium  sponge, 
absorbs  many  times  its  volume  of  hydrogen,  retaining  the  most  of  the  hydrogen  even 
at   100°.     At  a  high  heat  the  hydrogen  is  all  driven  off.     It  is  much  used  in  gas 
analysis  for  the  separation  of  hydrogen   from  other  gases  (Hempel,  B.,  1879,  12, 
636,   1006).     Also  used  for  scale  graduations  of  the  best  scientific  instruments. 

2.  Occurrence. — It  is  a  never-failing  element  in  the  platinum  ores,  native  or 
alloyed  with  Pt ,  Au  or  Ag  . 

3.  Preparation. — The  obtaining  of  pure  palladium  involves  its  separation  from 
the    other    platinum    metals,    i.  e.,    platinum,    iridium,    osmium,    rhodium    and 
ruthenium.     The  student  is  referred  to  the  various  works  on  metallurgy;  also 
to  the  following:  Bunsen,  A.,  1868,  146,  265;  Wilm,  B.,  1885,  18,  2536;  and  Mylius 
and  Forster,  B.,  1892,  25,  665. 

4.  Oxides  and  Hydroxides. — Palladium  monoxide,  PdO  ,  is  the  most  stable  of 
the  oxides   of  Pd  .     It  is  formed  by  the  gentle  ignition  of  Pd(N03)2   or  the 
precipitation   of   PdCl2   with   Na2CO3  ,   forming  Pd(OH)2  ,   and   then   igniting. 
Palladic  oxide,  PdO2  ,  when  gently  ignited  loses  half  its  oxygen,  becoming  PdO  . 

5.  Solubilities. — a. — Metal. — It   is   slowly   dissolved   by   boiling   with   HC1   or 
H2SO4;  HN03   dissolves  it,  even  in  the  cold,  forming  Pd(NO3)a  .     It  is  more 
readily   soluble   in   nitrohydrochloric    acid,    forming   PdCl4  .     It   is   not   at   all 
attacked  by  H2S  .     An  alcoholic  solution  of  iodine  blackens  it,  and  when  fused 
with  KHSO4  it  becomes  the  sulphate  (distinction  from  platinum).     6. — Oxides. — 
Pd02   is   soluble  in  HC1  with   evolution   of   Cl ,   forming  PdCl2  .     Pd(OH)2   is 
readily  soluble  in  acids  forming  palladous  salts,     c. — Suits. — Palladic  chloride, 
PdCl4  ,  the  most  stable  of  the  palladic  salts  is  decomposed  by  boiling  with 
water  or  by  much  dilution  with  cold  water,  forming  PdCl2  .     It  forms  double 
chlorides  with  other  metals,  as  calcium  palladic  chloride,  CaPdCl0  ,  which  for 
the  most  part  are  stable,  and  soluble  in  water  and  alcohol.     Potassium  palladic 
chloride,  K2PdCle  ,  is  but  sparingly  soluble  in  water,  insoluble  in  alcohol;  par- 
tially decomposed  by  both  solvents. 

Palladous  chloride  is  readily  soluble  in  water  with  a  brownish-red  color;  with 
metallic  chlorides,  it  forms  double  chlorides,  as  potassium-palladous  chloride, 
K.PdCl4  ,  all  of  which  are  soluble  in  water. — Palladous  iodide  is  insoluble  in 
water,  alcohol  or  ether;  insoluble  in  dilute  hydrochloric  acid  or  hydrioclic  acid; 
slightly  soluble  by  iodides  and  by  chlorides". — Palladous  nitrate,  Pd(N03)2  ,  is 
soluble  in  water  with  free  nitric  acid;  the  solution  being  decomposed  by  dilu- 
tion, evaporation,  or  by  standing,  with  precipitation  of  variable  basic  nitrates. — 
Palladous  sulphate,  PdSO4  ,  dissolves  in  water,  but  decomposes  in  solution  on 
standing. 

6.  Reactions.— Palladous  chloride  is  precipitated  by  potassium  hydroxide  or 
sodium    hydroxide;    as    brown    basic    salt    or    as    brown    palladous    hydroxide, 
Pd(OH),  ,  soluble  in  excess  of  the  hot  reagents.     Ammonium  hydroxide  gives 
a  flesh-red  precipitate  of  palladio-diammonium  chloride,  (NH3)2PdCl2  .     The  flesh- 
red  precipitate  is  soluble  in  excess  of  the   ammonia,   and   from   this   solution 
reprecipitated   by   hydrochloric    acid,   with   a   yellow    color.     The    fixed    alkali 
carbonates    precipitate    the    hydroxide;    ammonium    carbonate    acts    like    the 
hydroxide. — Potassium  cyanide  precipitates  palladous  cyanide,  Pd(CN)2  ,  white, 
soluble   in    excess    of    the   reagent.     Phosphates    give    a    brown    precipitate. — 
Hydrosulphuric    acid    and    sulphides    precipitate    the    dark-brown    palladous 
sulphide,  PdS  ,  insoluble  in  the  ammonium  sulphides,  soluble  in  nitrohydro- 
chloric acid.     Potassium  iodide  precipitates  palladous  iodide,  PdI2  ,  black,  visible 
in  500,000  parts  of  the  solution,  with  the  slight  solubilities  stated  in   5c,   an 
important  separation  of  iodine  from  bromine.     In  very  dilute  solutions,  only  a 


iRiDWM.  §106, 7. 

color  is  produced,  or  the  precipitate  separates  after  warming.    At  a  red  heat, 
the  precipitate  is  decomposed. 

Palladous  nitrate  gives  most   of  the   above   reactions;   no  precipitate  with 
ammonia,  and  a  less  complete  precipitate  with  iodides. 

7.  Ignition. — Nearly  all  the  palladium  compounds  are  reduced  by  heat,  before 
the  blow-pipe,  to  a  "  sponge."     If  this  be  held  in  the  inner  flame  of  an  alcohol 
lamp,  it  absorbs  carbon  at  a  heat  below  redness;  if  then  removed  from  the 
flame,  it  glows  vividly  in  the  air,  till  the  carbon  is  all  burnt  away  (distinction 
from  platinum). 

8.  Detection. — Palladium  is  precipitated  with  the  second  group  metals  by  H2S, 
not  dissolved  by  (NH4)2SX  (separation  from  the  tin  group).     It  is  distinguished 
from  mercury  by  its  precipitation  as  a  cyanide  with  mercuric  cyanide.     It  is 
precipitated  from  quite  dilute  solutions  by  KI   (distinction  from  Bi  and  Cd) ; 
an  excess  of  the  KI  dissolves  the  black  palladous  iodide,  PdI2  ,  to  a  dark  brown 
solution.     KCNS  does  not  precipitate  palladium  salts,  not  even  after  the  addi- 
tion of  SO2  (separation  from  Cu).     The  addition  of  H,SO4  and  alcohol  separates 
lead  from  palladium.     The  presence  of  the  metal  should  be  further  confirmed 
by  reduction  and  study  of  the  properties  of  the  "  sponge  "  obtained. 

9.  Estimation. — (Jf)   As  metallic  palladium,  to  which  state  it  is  reduced  by 
mercuric   cyanide   or   potassium   formate,    and   ignition,    first  in    the    air   and 
then  in  hydrogen  gas.     (2)  As  K2PdClG  .     Evaporate  the  solution  of  palladic 
chloride  with  potassium  chloride  and  nitric  acid  to  dryness,  and  treat  the  mass 
when  cold  with  alcohol,  in  which  the   double  salt  is  insoluble.     Collect  on   a 
weighed  filter,  dry  at  100°,  and  weigh. 

10.  Oxidation. — Palladium  is  reduced  as  a  dark-colored  precipitate,  from  all 
compounds  in  solution,  by  sulphurous  acid,  stannous  chloride,  phosphorus,  and 
all  the  metals  which  precipitate   silver   (§59,   10).     Ferrous   sulphate   reduces 
palladium  from  its  nitrate,  not  from  its  chloride.     Alcohol,   at  boiling  heat, 
reduces  it;  oxalic  acid  does  not  (distinction  from  gold  §73,  6&). 


§107.  Iridium.     Ir  =  193.1 .     Usual  valence  three  and  four. 

1.  Properties.— Specific  gravity,  22.421  (Deville  and  Debray,  C.  r.,  1875,  81,  839). 
Melting  point  2350°  ?  (Cir.  B.  S.,  35,   1915).     When  reduced  by  hydrogen  it  is 
a  gray  powder,  which  by  pressing  and  igniting  at  a  white  heat  changes  to  a 
metallic  mass  capable  of  takir.g  a  polish.     It  is  used  mostly  as  an  alloy  with 
platinum,   forming  a  very   hard,   durable   material   for  standard  weights   and 
measures.     A  platinum-iridium  dish  containing  25  to  30  per  cent  iridium  is  not 
attacked  by  nitrohydrochloric  acid. 

2.  Occurrence. — Found  in  platinum  ores,  usually  as  an  alloy  with  platinum 
or  osmium. 

3.  Preparation. — The  platinum   residues   are  mixed  with   Pb   and   PbO    and 
heated  at  a  red  heat  for  one-half  hour,  then  treated  with  acids.     The  residue 
contains  the  iridium  as  osmium-iridium  or  platinum-iridium  with  other  plat- 
inum metals.     This  residue  is  mixed  with  NaCl  in  a  glass  tube  and  heated  to 
a  red  heat  in  a  current  of  chlorine.     Much  of  the  osmium  passes  over  as  the 
volatile  perosmic  acid,  and  is  condensed.     The  double  sodium  chlorides  of  Ir  , 
Os  ,  Rh ,  Pt ,  Pd  and  Ru  are  dissolved  in  water,  filtered  and,  when  boiling  hot, 
decomposed  by  H2S  .     The  iridium  is  reduced  from  the  tetrad  to  the  triad,  but 
is  not  precipitated  until  after  all  the  other  metals.     By  stopping  the  current  of 
H2S  just  as  the  brown  iridium  sulphide  begins  to  form,  a  complete  separation 
can  be  made  by  filtration.     By  recrystallization  the  pure  sodium  double  salt, 
6NaC1.2lrCl3  +  24H20  ,  is  obtained,  which  is  changed  to  the  tetrad  ammonium 
double  salt,  (NH4)2IrCl6  ,  by  the  addition  of  NH4C1  and  oxidation  with  chlorine 
(Wohler,   Pogg.,   1834,   31,   161).     This   upon  ignition   gives   the   pure   metal   as 
iridium  sponge.     Or,  the  double  sodium  salt  is  ignited  with  sodium  carbonate, 
exhausted  with  water  and  reduced  by  ignition  in  a  current  of  hydrogen,  leav- 
ing the  metal  as  a  fine  gray  powder.     (See  also  §106,  3). 

4.  Oxides  and  Hydroxides. — Iridium  forms  two  series  of  oxides  and  hydrox- 
ides, the  metal  acting  as  a  triad  and  tetrad  respectively.     IrO3  is  formed  by 


§108,  5.  OSMIUM.  135 

igniting  the  metal  in  the  air  at  a  bright  red  heat,  henoe  the  scaling  of  platinum 
dishes  which  contain  iridium.  The  hydroxide,  Ir(OH)4  ,  is  formed  by  boiling-  a 
solution  of  the  trichloride,  IrCls  ,  in  a  fixed  alkali  hydroxide  or' carbonate. 
Careful  addition  of  KOH  to  IrCl3  in  a  vessel  full  of  liquid  and  closed  to  exclude 
air  gives  Ir(OH)3  ,  easily  oxidi/ed  to  Ir(OH)4  (Clans,  J.  pr.,  1846,  39,  104). 

5.  Solubilities. — Freshly  precipitated  iridium  may  be  dissolved  in  nitrohydro- 
chloric  acid.     The  ignited  metal  is  insoluble  in  all  acids.     Its  proper  solvent  is 
chlorine.     Iridium  trichloride,  IrCl3  ,  is  soluble  in  water  and  forms  with  the 
alkali  chlorides  double  chlorides,  soluble  in  water,  insoluble  in  alcohol.     The 
tetrachloride  with  sodium  chloride,  Na2IrCl«  ,  is  formed  when  the  platinum 
residues  mixed  with  NaCl  are  heated  in  a  current  of  chlorine.     It  is  soluble  in 
water.     The  corresponding  ammonium  salt  may  be  formed   from  the  sodium 
salt  by  precipitation  from  the  concentrated  solution  with  NH4C1 ,  a  reddish- 
brown  precipitate,  soluble  in  20  parts  of  water  (Vauquelin,  A.  Ch.,  1806,  59,  150 
and  225).     The  potassium  double  salt  is  sparingly  soluble  in  water. 

6.  Reactions. — Fixed  alkali  hydroxides  or  carbonates  precipitate  from  toil- 
ing solutions  of  iridium  chloride,  IrCl3  or  IrCl4  ,  iridiuju  hydroxide,  Ir(OH)4  , 
dark  blue,  insoluble  in  all  acids  except  HC1 .     Potassium  nitrite  added  to  a  hot 
solution  of  iridium  salts  gives,  first  a  yellow  color  and  finally  a  yellow  precipi- 
tate, insoluble  in  water  or  acids.    Hydrogen  sulphide  reduces  IrCl4  to  IrCl3  , 
and  then  precipitates  the  trisulphide,  Ir2S,  ,  brown,  soluble  in  alkali  sulphides. 

7.  Ignition.— When  iridium  is  fused  with  potassium  acid  sulphate  it  is  oxid- 
ized, but  does  not  go  into  solution  (difference  from  rhodium,  §105,  7).     Ignition 
on  charcoal  reduces  all  iridium  compounds  to  the  metal.     Fusion  in  the  air 
with  sodium  hydroxide  or  with  sodium  nitrate  causes  oxidation  of  the  metal, 
the  iridiiim  oxide  formed  being  partially  soluble  in  the  fixed  alkali. 

8.  Detection.— See  3  and  6. 

9.  Estimation. — It  is  converted  into  the  oxide  by  igniting  with  KN03   and 
then  reduced  by  ignition  in  an  atmosphere  of  hydrogen. 

10.  Oxidation. — Formic   acid    (from   hot   solution),   zinc   and   H2S04    or  HC1 
reduce   iridium   compounds  to   the   metal.     SnCl2  ,    FeSO4    and   H2C2O4    reduce 
tetrad  iridium  to  triad,  but  do  not  further  reduce  (separation  from  gold,  §73, 
60,  h  and  &). 

§108.  Osmium.    Os  =  190.9.    Valence  two  to  eight. 

1.  Properties. — Specific  gravity,  22,477,  the  heaviest  of  all  bodies  (Deville  and 
Debray,   C.  r.,   1876,  82,   1076).     Melting  point  2700°?   (Cir.   B.  S.,  36,   1915). 
In  the  absence  of  air  it  may  be  heated  above  the  vaporization  point  of  Pt  without 
melting  or  oxidizing.     In  presence  of  air,  when  heated  a  little  above  the  melting 
point  of  Zn,  it  burns  to  the  volatile  poisonous  perosmic  acid,  OsO4.     In  com- 
pact form  it  is  very  hard,  cutting  glass,  and  possesses  a  metallic  lustre,  with  a 
bluish  color  resembling  Zn. 

2.  Occurrence. — Always  present  in  the  residues  of  the  platinum  ores,  m  com- 
bination with  iridium. 

3.  Preparation. — The  iridium  osmium  alloy  or  other  Os  containing  material 
is  finely  divided  and  distilled  in  a  current  of  chlorine  or  with  nitrohydrochloric 
acid,  the  osmium  passing  into  a  receiver  containing  KOH.     By  repeated  addi- 
tions of  HNO3  and  further  distillation,  the  osmium  may  all  be  driven  into  the 
receiver.     The  distillate  is  treated  with  HC1  and  Hg  and  the  amalgam  ignited 
in  a  current  of  hydrogen  (Berzelius,  Pogg.,  1829,  15,  208). 

4.  Oxides.— Osmium   forms  five   different   oxides,   OsO  ,   Os203  ,   OsO,  ,   OsO.,  , 
Os04  .     The    first   three   are    bases,   the    salts   of   which   have    been    but   little 
studied;  OsO3  forms  salts  with  bases,  and  OsO4  acts  rather  as  an  indifferent 
peroxide.     Perosmic   acid,    OsO4  ,   exists   as   white   glistening   needles,   melting 
under  100°,  sparingly  soluble  in  water,  its  solution  having  a  very  penetrating 
odor,  resembling  that  of  chlorine.     The  fumes  of  the  acid  are  very  poisonous, 
and   cause  inflammation   of   the   eyes.     H2S   is   recommended   as   an   antidote 
(Clauss,  A.,  1847,  63,  355). 

5.  Solubilities.— The  metal  in  compact  condition  is  not  at  all  attacked  by  any 
acid.     The  precipitated  metal  is  slowly  dissolved  by  nitrohydrochloric  or  funv 


136  TUNGSTEN.  §108, 6. 

ing  nitric  acid.    By  heating  the  metal  in  a  current  of  chlorine  a  mixture  of 
OsCl2  and  OsCl4  is  formed.     They  are  both  unstable. 

6.  Reactions. — Perosmic  acid,  OsO4  ,  when  boiled  with  alkalis,  is  reduced  to 
osmates,   as   K2OsO4  .     A   solution   of    perosmic   acid   decolors   indigo,    oxidizes 
alcohol  to  aldehyde,  and  liberates  iodine  from  potassium  iodide.     In  the  pres- 
ence of  a  strong  mineral  acid,  H2S  precipitates  osmium  sulphide,  OsS4  ,  brown- 
ish black  (Claus,  /.  pr.,  1860,  79,  28);  insoluble  in  alkali  hydroxides,  carbonates 
or  sulphides. 

7.  Ignition. — Osmium  when  heated  on  a  piece  of  platinum  foil  gives  an  in- 
tensely luminous  flame  of  short  duration.     By  holding  the  foil  in  the  reducing 
flame  and  then  again  in  the  oxidizing  flame,  the  luminosity  may  be  repeated. 
If  a  mixture  of  the  metal  or  of  the  sulphide  and  potassium  chloride  be  heated 
in  a  current  of  chlorine,  a  double  salt  of  potassium  osmic  chloride  is  formed, 
sparingly  soluble  in  cold  \vater,  more  readily  in  hot  water.     Alcohol  precipitates 
it  from  its  solutions  as  a  red  crystalline  powder. 

8.  Detection. — By  the  intensely  luminous  flame  when  ignited  on  a  platinum 
foil;  by  oxidation  and  distillation  as  perosmic  acid  and  identification  by  odor, 
action  on  indigo  and  on  potassium  iodide. 

9.  Estimation. — It  is  weighed  as  the  metal  (see  3). 

10.  Oxidation. — Os04  is  reduced  to  Os02  by  ferrous  sulphate.     Zn  and  many 
other  metals  in  presence  of  strong  acids  precipitate  the  metal.     The  metal  is 
also  obtained  from  all  osmium  compounds  by  ignition  in  a  current  of  hydrogen. 

§109.  Tungsten  (Wolframium).    W  =  184.    Valence  two  to  six. 

1.  Properties.  —Specific  gravity,    18.71-18.74    (Z.  Anorg,   1909,  1910).     Melting 
point,   3000°    (Cir.   B.   S.,   35,    1915).     A  tin-white  or  steel-gray  metal,   brittle, 
harder   than    agate.     That    precipitated   from   acid   solutions   is   a   velvet-black 
powder.     Non-magnetic.     Stable  in  the  air  at  ordinary  temperature;    burning 
at  a  high  temperature,  it  decomposes  steam  at  a  red  heat.     Dry  metallic  tungsten 
powder,  compressed  into  a  bar  in  a  hydraulic  press,  may,  by  repeated  heatings 
and  swaging  or  rolling,  be  converted  into  a  ductile  form,  and  drawn  into  wire 
(Met.  and  Cham.  Eng.,  3,  XII,  1914). 

2.  Occurrence. — Tungsten  does  not  occur  in  nature  in  large  amounts,  nor  is 
it    widely    disseminated.     The    most    common    tungsten    minerals    are    scheelite, 
(CaWO4),  and  wolframite   (FeWO4  and  MnWO4 ,  in  variable  proportions).      It 
never  occurs  native. 

3.  Preparation. — By  reduction  of  WO3  in  H  at  a  red  heat   (Zettnow,  Pogg., 
1860,    111,   16);  by  ignition   of   W03    and   Na  under   NaCl  .     Tungstic   acid   of 
commerce  is  prepared  by  igniting  for  several  hours:  100  parts  Na2CO3  ,  ignited; 
150  parts  finely  ground  wolframite;  and  15  parts  NaNO3  .     The  cooled  mass  is 
exhausted  with  water  and  the  filtrate  poured  into  hot,  moderately  concentrated 
HC1  (Franz,  J.  pr.,  1871,  (2),  4,  238). 

4.  Oxides. — WO,  is  obtained  as  a  brown  powder  by  decomposing  WC14  with 
water  (Roscoe,  1.  c.).     WO3  is  a  lemon-3rellow,  soft  powder,  insoluble  in  water 
or  acids.     It  is  formed  by  ignition  of  the  metal,  lower  oxides  or  decomposable 
salts  in  the  air.     The  blue  tungsten  oxides  are  compounds  between  WO2  and 
W08. 

5.  Solubilities. — The  metal  is  scarcely  at  all  attacked  by  HC1  or  H2S04 ,  slowly 
by  HNO3  or  nitrohydrochloric,  slowly  soluble  in  alkalis.     The  halogens  com- 
bine directly  upon  heating.     WO2  is  readily  soluble  on  heating  with  HC1  and 
H2SO4    to   a  red   color.     It   is   also   soluble   in   KOH   with   red   color,    evolving 
hydrogen.     Both   the   acid   and    alkaline    solutions    deposit  the    blue    oxide   on 
standing  (von  der  Pfordten,  A.,  1884,  222,  158).     WO3  is  insoluble  in  water  or 
acids,  not  even  soluble  in  hot  concentrated  H2S04  .     Soluble  in  KOH  ,  K2CO, 
and  NH4OH  .     In  an  atmosphere  of  C02   it  reacts  with  the  chlorides  of  Ca  , 
Mg,  Co,  Ni  and  Fe  (not  with  those  of  Pb  ,  Ag  ,  K  and  Na),  e.g.,  MC12   + 
2W03  =  MW04  +  WO2C12  .     Heated  with  chlorine,  W02C12  is  formed,  and  also 
WC14  ,  decomposed  by  water.     S  ,  H2S   or  HgS  form  WS3  on  heating  with  W03  . 
Soluble  alkali  tungstates  are  formed  by  fusion  of  the  acid,  WO3  ,  with   the 
alkali  metal  carbonates,  more  slowly  by  boiling  with  the  carbonates.     Acids 
form,    from    solutions    of    the    alkali    tungstates,    a   white    precipitate    of    the 
hydrated  acid  turning  yellow  on  boiling,  insoluble  in  excess  of  the  acids  (dis- 


§111,    5.  GERMANIUM.  137 

tinction  from  MoO3),  soluble  in  NH4OH  .  Phosphoric  acid  changes  tungstic 
acid  to  the  metatungst  ic  acid,  which  is  soluble  in  water  arid  not  precipitated 
by  other  acids.  Long1  boiling1  of  the  solution  of  metatungsl  ic  acid  causes  1  lie- 
precipitation  of  tungstic  acid.  Fusion  of  W03  with  KHSO,  gives  a  compound 
of  potassium  tungstate  and  tungstic  acid,  uo1  readily  soluble  in  water  but  very 
readily  soluble  in  (NH4),CO3  (distinction  from  silica,  §249,  5). 

6.  Reactions. — Solutions  of  salts  of  Ba  ,  Ca  ,  Pb  ,  Ag  and  Hg  produce  white 
precipitates     with   solutions     of     alkali      tungstates.     H2S     precipitates      WS;. 
from   acid    solutions,    the    sulphide   dissolving   readily    in    (NH4)2S  ,   forming   a 
thiotungstate  (NH4)2WS4  .     The  tungstates,  like  the  molybdates,  form  complex 
compounds   with  phosphoric   acid,  i.e.,  phosphomolybdates   and   phospholuiig- 
states,  which  react  very  similarly  with  ammonium  salts  and  with  organic  bases 
(§75,    Grf).     K4Fe(CN)0   gives   with   tungstates    (in   presence   of   acids)    a   deep 
brownish-red  fluid,  forming  after  some  time  a  precipitate  of  the  same  color. 
Solution  of  taniiic  acid  gives  a  bro\vii  color  or  precipitate. 

7.  Ignition. — With  NaPO3  ,  W03  dissolves,  on  fusion,  to  a  clear  or  yellowish 
bead  in  the  oxidizing  flame;  in  the  reducing  flame  it  has  a  blue  color,  changing 
to  red  on  addition  of  FeSO4  .     Heated  on  charcoal  in  presence  of  Na2C03  with 
the  blow-pipe,  using  the  reducing  flame,  the  metal  is  obtained. 

8.  Detection. — If  a  tungstate  be  fused  with  Na2CO3  ,  the  mass  warmed  with 
water  and  the  water  then  absorbed  with  strips  of  filter  paper,  the  tungsten 
may  be  detected  by  moistening  the  strip  with  HC1  and  warming,  obtaining  the 
yellow  color  of  WO3;  and  the  blue  color  of  a  lower  oxide  by  moistening  with 
SnCl2  and  warming.     (NH4)2S  does  not  color  the  paper,  even  after  adding  HC1  , 
but  on  warming  a  blue  or  green  color  is  obtained. 

9.  Estimation. — It  is  converted  into  WO3  and  weighed  as  such  after  ignition. 

10.  Oxidation. — WO3  gives  with  SnCl,  ,  or  Zn  in  presence  of  HC1  or  H2SO4  , 
a  beautifid  blue  color,  due  to  the  formation  of  oxides  between  W02  and  W03  , 
blue  oxides  of  tungsten  (delicate  and  characteristic). 

§111    Germanium.      Ge  =  72.5.      Valence  two  and  four. 

1.  Properties.—  Specific  gravity,  5.469  at  20.4°  (Winkler,  /.  pr.,  1886,  (2),  34, 
177);    melting  point,  958°  (Cir.  B.  S.,  35,  1915).     A  gray-white  crystalline  metal. 
Fused  under  borax  it  gives  a  grayish-white  regulus  with  a   metallic  lustre.     It  is 
stable  in  the  air,  volatilized  at  a  high  heat  (Meyer,  B.,  1887,  20,  497),  and  is  easily 
pulverized.     It  burns  in  oxygen  to  form  germanic  oxide,  GeO2  . 

2.  Occurrence. — It  is  found  in   small   quantity  in   argyrodite,    a  sulphide  of 
silver    and    germanium,    (3Ag2S.GeS2)  ,    a    silver    ore    from     Freiburg,     Saxony; 
eux^ni  e    (msbate   and  titanate  of  yttrium,  erbium,  cerium  and  uranium)  from 
Sweden  (Kriiss,  C.  C.,  1888,  75);    also  contains  small  amounts  of  germanium. 

3.  Preparation.— It  is  formed  by  reduction  of  the  oxide,  GeO2  ,  with  H  ,  C 
or  Mg  (Winkler,  B.,  1891,  24,  891) ;  also  by  reduction  of  the  sulphide  in  H  . 

4.  Oxides. — It  forms  two  oxides,  GeO  and  GeO2  .     To  prepare  pure  Ge02  ,  the 
mineral  argyrodite  is  pulverized  and  intimately  mixed  with  equal  weights  of 
Na2C03  and  S  and  heated  to  a  good  full  ignition.     The  mass  must  be  added 
carefully  to  prevent  foaming.     The   fused  mass  is  exhausted  with  H2O  ,   the 
germanium  going  into  solution  as  a  thiosalt.     With  a  decided  excess  of  H2S04  , 
the  sulphide  is  completety  precipitated.     The  precipitate  is  now  dissolved  in 
KOH  ,  the  sulphides  of  A.g ,  Cu  and  Pb  remaining  undissolved.     By  adding  to 
the  KOH  solution  H2S04  not  quite  to  neutralization,  the  As  and  Sb  sulphides 
are   precipitated   on   boiling,   while   the   GeS   remains    in    solution   with    some 
As2S3;  H2S  is  carefully  added  to  the  solution  until  the  As2S3  is  all  precipitated, 
then  the  filtrate  is  made  strongly  acid  with  H2SO4  ,  and  the  solution  evaporated 
till  SO3  fumes  escape.     The  mass  is  dissolved  in  hot  water,  and  upon  cooling 
Ge02  crystallizes  out  (Winkler,  I.  c.). 

5.  Solubilities. — Germanium  is  insoluble  in  HC1 ,  soluble  in  nitrohydrochloric 
acid   as   GeCl4  ,   and   oxidized  with   HNO3    to  Ge02  .     Hot  concentrated   H2S04 
evolves  S02    and   forms   Ge(S04)2  .     Insoluble   in   KOH    solution    but   dissolves 
with   incandescence    in    fused   KOH .     It   unites   directly   with    01 ,   Br    and    I 
(Winkler,    I.  c.).     Germanic   oxide,    GeO2  ,   is    a  white   powder,   very    sparingly 
soluble  in  water  or  acids.     Fused  with  fixed  alkali  hydroxides  or  carbonates  it 
is  converted  into  compounds  soluble  in  water.    GeCl4  is  a  liquid,  boiling  at  84°; 


138  TELLURIUM.  §112,  6. 

it  is  decomposed  by  water.  If  a  solution  of  the  oxide  in  excess  of  HC1  be 
evaporated  to  dryness  the  Ge  is  all  volatilized.  GeS2  is  soluble  in  222  parts 
water,  in  alkali  sulphides  and  hydroxides;  insoluble  in  HC1  or  H2SO4  ,  which 
precipitate  it  from  its  solutions;  soluble  in  nitrohydrochlorie  acid  with  separa- 
tion of  sulphur.  Nitric  oxide  changes  it  to  Ge02  with  separation  of  sulphur. 

6.  Reactions. — Germanium  salts  give  almost  no  characteristic  reactions  with 
the  various  reagents.     H2S  precipitates  germanic  sulphide,  GeS2  ,  white,  from 
solutions  of  the  salts  quite  strongly  acid.     The  sulphide  is  soluble  in  ammonium 
sulphide,  forming  a  thio  salt,  thus  placing  Ge  in  division  A  of  the  second  group. 

7.  Ignition. — Heated  before  the  blow-pipe  in  the  reducing  flame  without  an 
alkaline  flux  the  metal  is  formed,  and  at  the  same  time  a  Avhite  coating  of 
the  oxide.     It  forms  a  colorless  bead  with  borax. 

8.  Detection. — In  the  mineral,  argyrodite,   by  heating  in  an   atmosphere  of 
H2S  or  illuminating  gas,  an  orange-yellow  sublimate  is  obtained,  which  may  be 
examined  under  the  microscope  and  in  the  wet  way   (Haushofer,  C.  C.,  1888, 
867). 

9.  Estimation. — It  is  converted   into  the   sulphide,   GeS2  ,   and   then   heated 
with  HNO3  and  weighed  as  GeO2  . 

10.  Oxidation.— Zn  in  acid  solutions  of  Ge  salts  precipitates  the  metal  as  a 
dark  brown  slime.     If  GeS-,  is  heated  in  a  current  of  H  ,  GeS  is  at  first  formed 
with  H2S,   finally  Ge°. 


§112.  Tellurium.     Te  =  127.5.      Valence  two,  four  and  possibly  six. 

1.  Properties. — Specific  gravity,  6.2445   (Berzelius,  Pogg.,  1834,  32,  1  and  577). 
Melting  point,  452°  (Cir.  B.  S.,  35,  1915).     Te  is  crystalline,  silver  white,  brittle, 
stable  in  the  air  and  in  boiling  water;    heated  in  the  air,  it  burns  with  a  greenish 
flame.     In  its  general  properties  and  reactions  it  stands  closely  related  to  S  and 
Se  (2). 

2.  Occurrence. — In  few  places  and  in  small  quantities  in  Germany,  Mexico, 
Bolivia,  United  States  and  Japan.     Some  of  the  minerals  are:    tellurite,  (TeO2)  ; 
tetradymite     (Bi2(TeS)3)  ;      ferrotellurite,      (FeTeO.,)  ,      sylvanite,     (AgAuTe4)  ; 
calaverite,  (AuTe2).     It  also  occurs  native. 

n.  Preparation. —  (1)  Fusion  with  alkali  carbonate  and  C  ,  which  converts  it 
into  a  telluride,  as  Na2Te;  then  solution  in  (air  free)  water,  the  air  being 
excluded  as  much  as  possible,  and  the  filtrate  precipitated  by  passing  air 
through  the  solution.  The  Te  is  precipitated  as  a  gray  metallic  powder,  con- 
taining what  Se  may  have  been  present.  (2)  Conversion  into  TeCl4  by  distilla- 
tion in  a  current  of  chlorine,  decomposition  of  the  chloride  with  water  to 
H2TeO3  and  precipitation  of  the  Te  with  KHSO3  .  (,?)  From  lead  chamber 
scale  by  digestion  with  Na2C03  and  KCN  ,  forming  KCNTe  .  The  decanted 
solution  is  acidified  with  HN03  and  the  Te  precipitated  with  H2S  (Schimose, 
C.  N.,  1884,  49,  157).  (4)  For  purification  of  the  commercial  Te  ,  see  Brauner 
(If.,  1889,  10,  411)  and  Schimose  (C.  N.,  1884,  49,  26,  and  1885,  51,  199). 

4.  Oxides  and  Hydroxides. — TeO  is  said  to  be  formed  by  heating  TeS03  in  a 
vacuum  above  180°:  TeSO.,  =  TeO  +  S02   (Divers  and  Schimose,  C.  N.,  1*8::,  47, 
221).     TeO2  forms  when  Te  is  burned  in  the  air,  and  when  TeCl4  is  decomposed 
by  boiling  water.     It  is  a  white  crystalline  solid,  sparingly  soluble  in   H2O  , 
more  soluble  in  acids  from  which  solutions  water  causes  a  white  precipitate  of 
Te02  or  H2TeO.,  .     H2TeO3  is  formed  when  a  HNO3  solution  of  Te  is  immediately 
poured  into  cold  water,  warming  to  40°  changes  it  to  TeO2  .     H2TeO4  is  made 
by  fusing  TeO,  with  KN03  ,  treating  the  K2Te04  so  *btained  with  soluble  lead 
or  barium  salt  and  decomposing  this  salt  with  H2SO4  or  H2S  ,  colorless  crystals, 
insoluble    in    alcohol    or    ether-alcohol    (separation    from    H2SO4).     It    can    be 
recrystallized  from  water  and  upon  heating  forms  TeO3   (Clarke,  Am.  8.,  1877, 
114,  281;  1878,  116,  401). 

5.  Solubilities. — Te  is  insoluble  in  HC1;  HN03   and  nitrohydrochloric  acids 
oxidize  it  to  H,TeO4;  in  H2SO4  it  becomes  H2Te03  with  evolution  of  SO,  (Hilger, 
A.,  1874.  171,  211):  soluble  in  warm  concentrated  solution  of  KCN,  from  which 
solution  HC1  precipitates  all  the  Te  .     H2TeO3  is  fairly  soluble   in  water,   red- 
c|ens  moist  litmus  paper  and  easily  decomposes  into  TeO2  and  H20  .    Acid  solu- 


**!."    rwifl 

°*   PHARMACY 

§113,  *.  SELENIUM.  139 

tions  of  Te02  are  precipitated  upon  addition  of  water  or  upon  standing.  Te02 
and  ILTeOj,  form  soluble  alkali  salts  with  «:lis  alkalis  from  which  solutions  of 
the  other  im-laJlic  sails  precipitate  the  respective  teliUrites.  H2Te04  is  soluble 
in  water,  acids  and  alkalis;  alkali  carbonates  form  acid  teimratoc,  less  soluble 
than  the  corresponding  normal  salts.  Solutions  of  the  alkali  tellurates  loiixi 
insoluble  tellurates  with  soluble  salts  of  the  other  metals,  e.  a.  K,TeO  4- 
BaCL  =  BaTe04  +  2KC1  . 

G.  Reactions. — Tellurium  is  classed  with  second  group  metals  because  of  its 
precipitation  from  solutions  of  tellurites  and  tellurates  by  H,S  .  The  precipi- 
tate is  not  a  sulphide,  but  is  Te  mixed  with  varying  proportions  of  S  ,  for  CS2 
removes  nearly  all  the  sulphur  (Becker,  A.,  1876,  180,  257).  In  appearance  the 
precipitate  of  Te  with  H.>S  very  much  resembles  SnS  ,  and  is  very  soluble  in 
(NH4)aS. 

At  a  high  temperature  Te  and  H  unite  directly,  forming  H2Te  (Brauner,  M., 
1889,  10,  446).  H2Te  is  best  prepared  by  heating  together  Te  and  Fe  or  Zn  and 
decomposing  these  tellurides  with  HC1  (analogous  to  the  corresponding  reac- 
tions with  sulphur,  §257,  4).  A  colorless  gas,  odor  similar  to  H2SJ  ,  burns  with 
a  blue  name,  fairly  soluble  in  water  and  is  precipitated  as  Te°  from  its  solution 
by  the  oxjrgen  of  the  air.  H2Te  precipitates  solutions  of  metallic  salts  very 
similarly  to  H2S  and  H2Se  . 

7.  Ignition. — Te  combines  on  ignition  with  most  metals  to  form  tellurides. 
TeO3  ignited,  decomposes  into  Te02  and  O  .     All  lower  Te  compounds  ignited 
with  KNO3  give  K2TeO4  .     All  Te  compounds  give  on  charcoal  with  the  blow- 
pipe a  white  powder,  which  colors  the  reduction  flame  green  and  disappears. 
Heated  in  an  open  glass  tube,  Te  compounds  give  a  sublimate  of  TeO2  .  which 
melts  upon  heating.     Te  compounds  fused  with  KCN  in  a  current  of  hydrogen 
form  potassium  tellurocyanate,  KCNTe;  soluble  in  water  but  precipitated  by  a 
current  of  air  as  Te°  (distinction  and  separation  from  Se).     Heated  with  Na2C03 
on  charcoal  Te  compounds  give  Na,Te  ,  which  blackens  silver  with  formation 
of  Ag2Te  . 

8.  Detection. — By  reduction  to  Te°  and  solution  in  cold  concentrated  H2S04 
to  a  purplish-red  solution  (characteristic).     Separated  from  Se  by  fusion  with 
KCN  in  a  current  of  hydrogen  and  precipitation  from  the  solution  by  a  current 
of  air. 

9.  Estimation. — The  Te  compound  is  heated  in  a  current  of  Cl  ,  TeCl4  being 
sublimed.     This  is  decomposed  by  water  to  Te02  ,  which  is  reduced  to  Te°  by 
S02  and  weighed  as  such  after  drying  at  100°  . 

10.  Oxidation. — Hydrogen  at  a  high  temperature  reduces  Te  compounds  to 
H2Te  .     H,S  reduces  Te  compounds  to  Te°  mixed  with  S  .     Fusion  with  KNO3 
oxidizes  all  Te  compounds  to   K2TeO4  .     S02   reduces  Te   compounds  to   Te°  . 
SnCL  and  Zn  in  acid  solutions  give  with  Te  compounds  a  black  precipitate 
of  Te°  .     Te  compounds  warmed  with  dextrose  in  alkaline  solution  are  reduced 
to  Te°  .     Tellurates  boiled  with  HC1  evolve  chlorine  and  are  reduced  to  H2TeO3  , 
which  precipitates  as  Te02  on  adding  water  if  too  much  HC1  be  not  present 
(distinction  from  Se). 


§113.  Selenium.     Se  =  79.2  .     Valence  two  and  four,  possibly  six. 

1.  Properties. — Specific  gravity,  red,  cryst.,  4.47;    gray  metallic  4.80;  («/.  phys. 
Chem.  4,  491;    1900).     Melting  point  217-220°  (Cir.  B.  S.,  35,  1915).     The  molten 
Se  does  not  become  completely  solid  until  cooled  to  50°.     Selenium  with  tellurium 
is  closely  related  to  sulphur,  and  like  sulphur  exists  in  amorphous  forms  (§266,  1). 
The    precipitated    Se   is    red.     The    brown    or   brown-black  powder  obtained  by 
quickly  cooling  from  the  molten  state  is  insoluble  in  €82.      Boiling  point,  676° 
to  683°  (Carnelley  and  Williams,  C.  N.,  1879,  39,  286). 

2.  Occurrence. — In  no  place  abundantly;  never  native.     It  is  found  in  com- 
bination with  minerals  in  the  Hartz  Mountains,  Sweden,  Argentine  Republic  and 
Mexico  (Billandot,  C.  N.,  1882,  46,  60).     It  occurs  in  very  small  quantities  with 
some  sulphides  of  Fe  ,  Cu  and  Zn  . 


140  SELENIUM.  §113,  3. 

3.  Preparation. — In  the  lead  chambers  of  the  H2S04  works  it  is  found  as  a 
red  deposit  with  some  S  ,  As2O3  ,  Sb,O3  ,  PbSO4  ,  etc.     The  scale  is  washed  with 
water  and  digested  with  KCN  solution  at  80°  to  100°,  until  the  red  color  entirely 
disappears.     The  filtrate    is    then   treated   with   HC1  ,  which   precipitates  the   £e  . 
It    is    further    purified   by  oxidation  to   SeO9 .    sublimed    and   then    reduced    with 
S02  (Nilson,  B.,  1874,  7,  1719). 

4.  Oxides  and  Hydroxides. — H2SeO3  is  prepared  by  oxidizing  Se  with  HNO3  , 
or  nitrohydrochloric  acid.     H,Se03  evaporated  to  dryness  gives  H2O  and  Se02  , 
crystalline.     SeO,   is  also  formed  by  burning  Se  in  air  or  oxygen;  it  has  an 
odor  similar  to  decaying  radish.     It  sublimes  at  250°-280°   as  a    yellow  vapor, 
condensing  to  white  needles  on  cooling.     SeO3  is  not  known.     H.Se04  ,  pure, 
is  a  white  crystalline  mass,  melting  at  58°.     H2Se04.H,0  is  crystalline  at  — 38°, 
and  if  recrystallized  melts  at  25°.     The  selenic  acid  usually  obtained  is  a  thick 
oily  liquid,  resembling  H2S04  and  containing  about  95  per  cent  H2Se04  .     It  is 
obtained  by  fusing  Se  or  SeO2  with  KN03  and  precipitation  of  the  K,SeO4  with 
soluble  salts  of  Ba  ,  Pb  ,  Ca  or  Cu  and  decomposing  the  washed  precipitates, 
suspended  in  water,  with  H2S04  or  H2S  . 

5.  Solubilities. — Se  dissolves  in  cold  concentrated  H2S04  to  a  green  colored 
solution  without  oxidation    (dilution  with  water  precipitates   the   Se) ;   if   the 
solution  be  warmed  SO,   is  evolved  and  the  green  color  disappears    (dilution 
with  water  gives  precipitate),  the  Se  being  oxidized  to  SeO,  .     HNO3  and  nitro- 
hydrochloric acid  oxidize  it  to  SeO2  .     Selenous  oxide,  SeO2  ,  is  soluble  in  water 
in  all  proportions,  forming  H2Se03  .     The  selenites  and  selenates  of  the  alkaline 
earths  are  insoluble  and  may  be  formed  by  adding  a  solution  of  the  metal  to 
an  alkali  selenite  or  selenate,  e.  g.,  Na2Se03  +  BaCl2  =  BaSe03  +  2NaCl .     Many 
of  the  selenites   are   soluble  in  excess  of  H2Se03  .     Selenates   are   less   stable 
than   selenites.     BaSeO4    is   soluble   in    HC1    (distinction    and    separation    from 
BaS04)  and  upon  long-continued  boiling  is  reduced  to  BaSeO,  . 

6.  Reactions. — Selenous  acid  precipitates  with  H2S  a  mixture  of  Se  and  S  , 
lemon  yellow,  bright  red  upon  heating   (Divers  and  Shimose,  C.  N.,  1885,   51, 
199).     This  mixture  is  soluble  in  (NH4)2S  ,  hence  in  qualitative  analysis  Se  is 
classed  among  the  metals  of  division   A,  second  group,   while  because   of  its 
general  properties  it  belongs  with  sulphur.     "When  Se  and  H  are   heated  to- 
gether they  begin  to  combine  directly  at  250°,  forming  H_,Se  (Ditte,  C.  r,,  1872, 
74,  980);  which  in  practically  all  its  reactions  is  similar  to  H^S  .     H.,Se  is  also 
formed  by  treating  K2Se  ,  FeSe  ,  etc.,  with  dilute  HC1  or  H3SO«;  HN03  gives 
H,SeO3  with  selenides.     H2Se  is  a  colorless  gas,  odor  similar  to  HoS  but  more 
penetrating.     It  is  more  poisonous  than  H2S  ,  burns  when  ignited,  combines 
slowly  but  completely  with  Hg°  ,    evolving  hydrogen.     100  cc.  of  water  dissolves 
331  cc.  of  the  gas  at  13°,  the  solution  reacting  acid  and  depositing  red  flakes  of 
Se  on  standing.     It  precipitates  the  selenides  of  the  metals  having  almost  the  same 
solubilities  as  the  corresponding  sulphides   (von  Reeb,  J.   Pharm.,   1869,    (4),  9 
173).     With  soluble  sulphites  H2Se  gives  a  precipitate  of  a  mixture  of  Se  and  S'. 

7.  Ignition. — When  Se  or  compounds  of  Se  are  fused  with  KCN  in  a  current 
of  hydrogen,  potassium  selenocyanate,  KCNSe  ,  is  formed.     Long  boiling  with 
HC1  separates  the  Se  ,  but  this  does  not  take  place  on  exposure  of  the  solution 
to  the  air  (separation  from  tellurium).     Selenium  compounds  heated  on  char- 
coal with  Na2C03  are  changed  to  Na2Se  ,  which  yields  a  black  stain  with  Ag° 
and  H2Se  with  dilute  acids. 

8.  Detection. — If  in  solution  as  selenites  it  is  precipitated  with  H2S  (soluble 
in  (NH4)2S);  oxidized  to  SeO,  and  obtained  as  the  white  needles  by  sublima- 
tion, and  reduced  from  its  solution  in  water  to  the  red  Se°  by  S02  .     If  present 
as  selenides,  decomposed  by  HC1  or  H2SO4  ,  forming  H2Se  ,  which  is  conducted 
into  water  and  the  Se°  precipitated  by  passing  air  or  oxygen  through  the  solu- 
tion. 

9.  Estimation. — Oxidized    to    selenic    acid    and   precipitated    as    BaSe04    and 
weighed  as  such.     If  BaS04  be  present  the  precipitate  is  reduced  in  H  ,  and 
the  resulting  BaSeO.,  separated  by  solution  in  HC1  .     Selenides  are  heated  in  a 
current  of  chlorine  in  a  hard  glass  tube,  being  converted  into  SeCl4  ,  which 
vaporizes  and  is  decomposed   in  water;   continued  chlorination   of  the  water 
solution  forms  H2SeO4  . 

10.  Oxidation.— Se°  is  oxidized  to  SeO2  by  HN08  ,  nitrohydrochloric  acid, 


§114.  THE  IRON  AND  ZINC  GROUPS.  141 

H2SO4  hot  concentrated,  by  heating  in  air  or  oxygen,  etc.  H2SeO3  is  oxidized 
to  H2SeO4  by  continued  chlorination,  and  by  fusion  with  KN03  .  H2SeO4  is 
reduced  to  H2Se03  by  boiling-  with  HC1 .  SO,  reduces  selenous  compounds  to 
the  red  Se°  ,  even  in  H2SO4  solutions  (distinction  from  tellurium)  (Keller, 
J.  Am.  Soc.,  1900,  22,  241)."  H2S  forms  a  precipitate  of  Se  mixed  with  S  .  SnCl2 
precipitates  Se°  from  HC1  or  H2SO4  solutions  of  selenous  compounds. 


THE  IRON  AND  ZINC  GROUPS  (THIRD  AND  FOURTH  GROUPS). 

114.  The  Metals  of  the  Earths  and  the  more  Electro-Positive  of  the 

Heavy  Metals. 


Aluminum.,, Al    =    27.1 

Chromium Cr    =     52.0 

Iron  Fe  '=     55.84 

Cobalt Co    =     58.97 

Nickel Ni    =     58.68 

Manganese Mn  =     54.93 

Zinc Zn  =     65.37 

Cerium Ce    =  140.25 

Columbium Cb   =    93.5 

Erbium E      =  167.7 

Gallium Ga  =     69.9 

Glucinum Gl    =       9.1 

Indium In    =  114.8 

Lanthanum La    =  139.0 


Neodymium Nd  =  144.3 

Praseodymium Pr  =  140.9 

Samarium Sa  =  150.4 

Scandium Sc  =     44.1 

Tantalum Ta  =  181.5 

Terbium Tb  =  159.2 

Thallium .Tl  =  204.00 

Thorium Th  =  232.4 

Titanium Ti  =     48.10 

Uranium U  =  238.2 

Vanadium V  =    51 .0 

Ytterbium Yb  =  173.5 

Yttrium Y  =  188.7 

Zirconium Zr  =    90.6 


§115.  The  metals  above  named  gradually  oxidize  at  their  surfaces  in 
the  air,  and  their  oxides  are  not  decomposed  by  heat  alone.  Zinc,  iron, 
cobalt,  nickel,  and,  with  more  difficulty,  manganese,  chromium,  and  most 
of  the  other  metals  of  the  groups,  are  reduced  from  their  oxides  by  igni- 
tion at  white  heat  with  charcoal.  They  are  all  reduced  from  oxides  by 
the  alkali  metals.  Iron  is  gradually  changed  from  ferrous  to  ferric 
combinations  by  contact  with  the  air.  Chromium  and  manganese  are 
oxidized  from  bases  to  acid  radicals  by  ignition  with  an  active  supply  of 
oxygen  in  presence  of  alkalis;  these  acid  radicals  acting  as  strong  oxidizing 
agents  (§8,  §9). 

§116.  The  oxides  and  hydroxides  of  these  metals  are  insoluble  in  water 
and  they  are  precipitated  from  all  their  salts  by  alkalis.  In  the  case  of 
zinc,  the  precipitate  redissolves  in  all  the  alkalis;  the  aluminum  hydroxide 
redissolves  in  the  fixed  alkalis,  but  very  slightly  in  ammonium  hydroxide; 
the  precipitate  of  chromium  redissolves  in  cold  solution  of  fixed  alkalis, 
precipitating  again  on  dilution  and  boiling;  the  hydroxides  of  cobalt  and 
nickel  dissolve  in  ammonium  hydroxide.  The  oxide  of  chromium  after 
ignition  is  insoluble  in  acids;  the  oxides  of  aluminum  and  iron  are  soluble 
with  difficulty. 

The  presence  of  tartaric  acid,  citric  acid,  sugar,  and  some  other  organic 
substances,  prevents  the  precipitation  of  bases  of  these  groups  by  alkalis. 


142  THE  IRON  AND  ZINC  GROUPS.  §117. 

§117.  Ammonium  salts,  as  NH4C1 ,  dissolve  moderate  quantities  of  the 
hydroxides  of  manganese,  zinc,  cobalt,  nickel,  and  ferrous  hydroxide;  but, 
so  far  from,  dissolving  the  hydroxide  of  aluminum,  they  lessen  its  slight 
solubility  in  ammonium  hydroxide. 

§118.  It  thus  appears  that  ammonium  hydroxide,  with  ammonium 
chloride,  the  latter  necessary  on  account  of  magnesium  (§189,  6a),  man- 
ganese (§134,  6a),  and  aluminum,  will  fully  precipitate  only  aluminum, 
chromium,  and  iron  of  the  important  metals  above  named.  These 
metals  therefore  constitute  the  THIRD  GROUP  (§127),  and  the  re- 
agent of  this  group  is  AMMONIUM  HYDROXIDE  in  the  presence  of 
AMMONIUM  CHLORIDE.  Since  aluminum,  chromium,  and  iron  are 
precipitated  by  ammonium  hydroxide  in  the  presence  of  ammonium 
chloride  (Fe"  by  its  previous  oxidation  with  HN03  is  present  as  Fe'") 
constituting  the  THIRD  GROUP;  the  remaining  of  the  most  important 
metals — cobalt,  nickel,  manganese,  and  zinc — constitute  the  FOURTH 
GROUP  (§137).  They  are  precipitated  by  the  group  reagent,  AMMON- 
IUM SULPHIDE  or  HYDROSULPHURIC  ACID  in  an  AMMONIACAL 
SOLUTION.  Some  chemists  do  not  make  this  classification  of  these 
metals,  but  precipitate  them  all  as  one  group  with  ammonium  sulphide 
(§144),  from  neutral  or  ammoniacal  solutions.  The  sulphides  of  Fe,  Co, 
Ni ,  Mn  ,  and  Zn  are  not  formed  in  presence  of  dilute  acids,  which  acids  keep 
them  in  solution  during  the  second  group  precipitation;  but  are  insoluble 
.in  water,  which  enables  them  to  be  precipitated  by  alkali  sulphides,  and 
separated  from  the  fifth  and  sixth  groups.  The  other  two  metals,  Al  and 
Cr,  do  not  form  sulphides,  in  the  wet  way,  but  are  precipitated  as  hy- 
droxides by  the  alkali  sulphides. 

§119.  Hydrosulphuric  acid  scarcely  precipitates  the  metals  of  these 
groups,  unless  it  be  from  some  of  their  acetates  (§135,  6e),  owing  to  the 
solubility  of  the  sulphides  in  the  acids,  which  would  be  set  free  in  their 
formation.  Thus,  this  change  cannot  occur — FeCl2  +  H2S  =  FeS  -|- 
2HC1 — because  the  two  products  would  decompose  each  other.  Hydrogen 
sulphide  does  not  precipitate  the  metals  of  these  groups  in  acid  solution 
unless  the  acid  is  very  weak  (acetic  acid  §135,  6e).  The  hydrogen  ions  of 
strong  acids,  which  are  largely  dissociated,  reduce  the  concentration  of  the 
sulphur  ion  of  the  hydrogen  sulphide  below  the  point  where  it  can  pre- 
cipitate the  sulphides  of  these  metals.  For  the  same  reason  these  sul- 
phides are  dissolved  by  strong  acids  and  the  reaction  FeCL  +  H2S  <=> 
FeS  -f  2HC1  cannot  proceed  from  right  to  left. 

As  acetic  acid  and  other  weak  acids  are  only  slightly  dissociated,  the 
concentration  of  the  hydrogen  ion  is  very  nmcli  less  and  the  decrease  in  the 
concentration  of  the  sulphur  ion  of  the  hydrogen  sulphide  is  slight.  The 
soluble  sulphides  axe  dissociate^  to  a  muqh  greater  extent  giving  a  con- 


THE  IRON  AND   ZINC  GROUPS.  143 

centration  of  the  sulphur  ion  sufficient  to  precipitate  the  sulphides  of  these 
metals.  (See  45.)  Therefore  when  ii  is  desired  i<>  precipitate  the  metals  as 
sulphides,  neutralized  hydrosulphuric  acid — an  alkali  suljihitle — is  used  in 
neutral  or  alkaline  solution;  or,  what  is  equivalent,  hydrosulphuric  acid 
gas  is  passed  into  the  strongly  ammonifical  xolnlion. 

§120.  As  most  of  the  normal  chemically  salts  of  heavy  metals  are  hydro- 
lyzed,  in  water,  giving  free  acids,  so  that  their  solutions  have  an  acid  reaction 
to  test-paper,  we  can  onlv  assure  ourselves  of  the  requisite  neutrality  by 
adding  sulTieient  ammonium  hydroxide,  which,  itself  precipitates  the  larger 
number  of  the  bases,  as  we  have  just  seen  (§116).  But  the  resulting. 
precipitate  of  hydroxide,  as  Fe(OH)2,  is  immediately  changed  to  sul- 
phide, FeS ,  by  subsequent  addition  of  ammonium  sulphide;  as  the  student 
may  observe,  by  the  change  in  the  color  of  the  precipitate. 

Ferric  and  manganic  salts  are  reduced  to  ferrous  and  mauganous  salts, 
by  hydrosulphuric  acid,  in  solution,  with  a  precipitation  of  sulphur,  and 
the  corresponding  reaction  occurs  with  chromates. 

§121.  Soluble  carbonates  precipitate  all  the  metals  of  these  groups,  in 
accordance  with  the  general  statement  for  bases  not  alkali  (§205,  6a). 
With  aluminum  and  chromium,  the  precipitates  dissolve  sparingly  in  ex- 
cess of  potassium  or  sodium  carbonate;  with  Co,  Ni  and  Zn  ,  the  precipitate 
dissolves  in  excess  of  (NH4)2C03 .  In  the  case  of  ferrous  and  manganous 
salts,  the  precipitates  are  normal  carbonates;  with  zinc,  cobalt,  and  nickel 
salts,  they  are  basic  carbonates;  while  with  ferric,  aluminum,  and  chrom- 
ium salts,  the  precipitates  are  hydroxides.  Barium  carbonate  precipitates 
Al ,  Cr'"  and  Fe'",  which,  in  the  cold  and  from  salts  not  sulphates,  is  a 
separation  from  the  fourth  group  metals. 

§122.  Soluble  phosphates  precipitate  these  as  they  do  other  non-alkali 
bases.  The  acid  solutions  of  phosphates  of  the  metals  of  the  third  and 
fourth  groups  are  precipitated  by  neutralization.  Phosphates  of  Co ,  Ni , 
and  Zn  are  redissolved  by  excess  of  NH4OH ,  and  those  of  Al ,  Cr ,  and  Zn 
by  excess  of  the  fixed  alkalis.  The  recently  precipitated  phosphates  of  all 
the  metals  of  these  groups  which  form  sulphides,  are  transformed  to  sul- 
phides by  ammonium  sulphide,  due  to  the  fact  that  the  sulphide  is  less 
soluble  than  the  phosphate:  FeHP04  -f  (NH4)2S  =  FeS  +  (NH4)2HP04 . 
Hsnce,  the  only  phosphates  which  may  occur  in  a  sulphide  precipitate  are 
those  of  Al ,  Cr ,  Ba  ,  Sr ,  Ca  ,  and  Mg  . 

§123.  The  rnetals  of  the  third  and  fourth  groups  are  not  easily  reduced 
from  their  compounds  to  the  metallic  state  by  ignition  before  the  blow- 
pipe, even  on  charcoal,  except  zinc,  which  then  vaporizes.  Three  of  them, 
however — iron,  cobalt,  and  nickel — are  reducible  to  magnetic  oxides.  The 
larger  number  of  them  give  characteristic  colors  to  beads  of  borax  and  of 
microcosmic  salt,  fused  on  a  loop  of  platinum  wire  before  the  blow-pipe. 


144  ALUMINUM.  §124,  1. 

None  of  them  color  the  flame  or  give  spectra,  unless  vaporized  by  a  higher 
temperature  than  that  of  a  Biinsen  burner  (spark  spectra). 

THE  IROX  GROUP   (THIRD  GROUP). 

Aluminum,  Chromium,  Iron. 
§124,  Aluminum.     Al  —  27.1.     Valence  three. 

1.  Properties.—  Specific  gravity,  2.708  (C.  JV.f  105,  1912).    The  cast  metal  has 
specific  gravity  of  2.56.     Melting  point,  658.7°   (Cir.   B.  S.,  35,   1915).     It  is  a 
tin-white  metal    (the   powder  is   gray),   odorless   and   tasteless,  very  ductile  and 
malleable,  about  as  hard  as  silver.     Its  boiling  point  is  above  2200°.     Impurities 
increase  the  melting  point.     When  molten  it  possesses  great  fluidity.     As  a  con- 
ductor of  heat  it  is  about  twice  as  good  as  tin  and  about  one-third  as  good  as 
silver.     It  conducts  electricity  about  one-half  as  well  as  copper  and  silver  (Dewar 
and  Fleming,  Phil.  Mag.,  (5)  36,  271,  1893.     Roy,  Inst.  Gt.  Brit.,  June  5,  1896), 
and  about  three  times  better  than  iron.     Commercial  aluminum  is   never   pure, 
containing  small  amounts  of  silicon  and  iron,  and  sometimes  Cu  and  Pb  ,  with 
96  to  99.75  per  cent  aluminum.     It  is  used  for  cooking  utensils,  canteens  and  other 
military  equipments,   boats,   small  weights,   measures,   articles  of  ornament   and 
scientific  instruments;    as  an  alloy  with  copper  (aluminum  bronze)  it  finds  exten- 
sive application. 

2.  Occurrence. — Not  found  free  in  nature.    Is  found  in  corundum,    ruby   anc1 
sapphire,    as    nearly    pure    A12O3  ;    in    diaspore      (A1(OH)3.A12O3)  ;    in    bauxite 
(ALOs.xHzO)  ;      in    orthoclase    (K£lSi,jOs)  ,      and     other    feldspars  ;   in    cryolite 
(Na^AlFe)  .     As  a  silicate  in  all  clays  and  in  very  many  minerals.     It  is  widely 
distributed,  constituting  about  one-twelfth  of  the  earth's  crust. 

3  Preparation. — (i)  By  electrolysis  of  the  fused  NaAlCl4  .  (2}  By  fusion  of 
cryolite  or  the  chloride  with  Na  or  K  .  (3)  By  heating  NaAlCl4  with  zinc,  with 
which  it  forms  an  alloy  from  which  the  zinc  is  driven  off  by  a  white  heat. 
(4)  By  fusion  of  the  chloride  with  potassium  cyanide.  (5}  By  fusing  ALS3 
with  iron.  A  great  many  new  methods  have  been  patented.  (6)  Aluminum 
is  prepared  commercially,  by  electrolysis  of  aluminum  oxide  dissolved  in  a  bath  of 
cryolite  (Na-iAlF6).  The  metal  is  deposited  around  the  cathode,  oxygen  being 
evolved  at  the  anode.  See  Dammer,  3,  79. 

4.  Oxide    and    Hydroxides. — A12O3    is    formed    by    heating-    the    hj-droxide, 
nitrate,  acetate  or  other  organic  salt,  difficultly  soluble  in  acids  after  ignition, 
but    may    be    dissolved    after    fusion    with    KHSO4    or    Na^CO.,  .     A1(OH)3    is 
formed  when  aluminum  salts  are  precipitated  with  cold  ammonium  hydroxide. 
A120(OH)4  is  formed  if  the  precipitation  is  made  at  100°. 

5.  Solubilities. — a. — Metal. — Pure  aluminum  scarcely  oxidizes  at  all  in  dry  or 
moist  artr;  the  electrolytically  deposited  powder  oxidizes  gradually  in  the  air. 
Powdered  or  leaf  aluminum  when  boiled  with  water  evolves  hydrogen,  forming 
the   hydroxide.     It   is  -attacked   by   the   halogens    forming   the    corresponding 
halides   (Gustavson,  BL,   1881,    (2),  36,  556).     Dilute   sulphuric  acid  attacks  it 
slowly,  evolving  hydrogen   (Ditte,  C.  r.,  1890,  11O,  573);  the  hot  concentrated 
acid   dissolves   it  readily  with   evolution   of   S02  .     Nitric   acid,   dilute   or   con- 
centrated, attacks  it  very  slowly  (Deville,  A.  CJi.,  1855,  (3),  43,  1-1:  Montemartini, 
Gazzetta,  1892,  22,  397;  Ditte,  I.e.,  782).     Hydrochloric  acid,  dilute  or  concen- 
trated, dissolves  it  readily  with  evolution  of  hydrogen;  also  attacked  readily 
by   fixed   alkalis,  sparingly   by   NH4OH    (Gottig,   J5.,   1896,   29,    1671),   evolving 
hydrogen  with  formation  of  an  aluminate:  2A1  +  2KOH  +  2H20  =  2KA10,  -f 
3H2  .     It  is  attacked  by  fixed  alkali  carbonates  (D.,  3,  87).     When  ignited  with 
sodium  carbonate,  aluminum  oxide  is  formed,  sodium  is  vaporized  and  a  small 
amount  of   aluminum   nitride  produced    (Mallet,  J.   C.,   1876,  30,   349).     Fused 
KOH  is  decomposed  by  aluminum  at  very  high  temperature,  the  potassium 
being  vaporized  (Deville,  J.,  1857,  152).     It  is  not  at  all  attacked  by  cold  four 
per  cent  acetic 'acid  (vinegar)  even  in  presence  of  NaCl ,  and  when  boiled  for 


§124,  6a.  ALUMINUM.  145 

14  hours  with  the  above  mixture  a  square  meter  of  surface  (weighing  24.7426  grams) 
lost  but  0.047  gram  (one  part  in  526). 

b. — Oxide  and  hydroxide. — The  oxide  is  insoluble  in  water,  and  when  not  too 
strongly  ignited  dissolves  readily  in  dilute  acids  and  in  fixed  alkalis.  Corundum, 
crystallized  A12O3 ,  is  insoluble  in  acids,  but  is  rendered  soluble  by  fusion  in  fixed 
alkali  carbonates  or  sulphates.  The  hydroxide  A1(OH)3  is  insoluble  in  water,  readily 
soluble  in  acids  and  in  fixed  alkalis,  sparingly  soluble  in  ammonium  hydroxide,  the 
solubility,  however,  being  much  decreased  by  the  presence  of  ammonium  salts. 
c. — Salts. — Aluminum  phosphate  is  insoluble  in  water.  The  normal  acetate  is 
soluble,  the  basic  acetate  insoluble  in  water  (separation  from  Cr  and  the  fourth 
group)..  The  chloride  is  deliquescent.  The  double  sulphates  of  aluminum  and  the 
alkali  metals  (alums)  are  soluble  and  readily  melt  in  their  water  of  crystallization, 
becoming  anhydrous.  Solutions  of  normal  salts  of  aluminum  have  an  acid  reaction. 

6.  Reactions,  a. — The  alkali  hydroxides  and  carbonates  *  precipitate 
aluminum  hydroxide  (1),  A1(OH)3  (4),  grayish -white,  gelatinous  insoluble 
in  water,  soluble  in  excess  of  the  fixed  alkali  hydroxides  f  "(£)  (Prescott, 
/.  Am.  Soc.,  1880,  2,  27;  Ditte,  A.  Ch.,  1897 '(6),  30,  266),  sparingly  soluble 
in  the  fixed  alkali  carbonates  and  in  ammonium  hydroxide  but  much  less 
so  if  ammonium  salts  be  present.  The  solution  of  fixed  alkali  alnminate 
is  precipitated  as  aluminum  hydroxide  by  careful  neutralization  of  the 
alkali  with  acids  including  hydrosulphuric  (3),  and  carbonic,  as  basic 
hydroxide,  by  adding  excess  of  ammonium  chloride  (4)  (distinction  from 
zinc  which  is  precipitated  by  a  small  amount  of  NH4C1 ,  but  redissolves  on 
adding  an  excess)  (Lowe,  Z.,  1865,  4,  350).  The  excess  of  potassium 
hydroxide  liberates  ammonia  forming  potassium  chloride,  thus  reducing 
the  amount  of  fixed  alkali  present.  The  precipitate  is  more  compact  and 
washes  more  readily  than  the  gelatinous  normal  hydroxide.  Barium  car- 
bonate, on  digestion  in  the  cold  for  some  time  completely  precipitates 
aluminum  salts  as  the  hydroxide  (5)  mixed  with  a  little  basic  salt.  (See 
§126,  6a.)  The  presence  of  citric,  oxalic,  or  tartaric  acid  greatly  hinders 
the  precipitation  of  aluminum  hydroxide,  and  an  excess  may  entirely  pre- 
vent its  precipitation  by  the  formation  of  a  soluble  double  salt,  e.  g., 
KA1(C4H406)L,  .  Other  organic  substances,  as  sugar,  pieces  of  filter  paper, 
etc.,  hinder  the  precipitation.  To  obtain  complete  precipitation  all  or- 
ganic substances  should  be  decomposed. 

(1)  A1C1,  +  3KOH  =-Al(OH),  +  3KC1 

2A1C13  +  3K2CO,  +  3H20  =  2A1(OH)3  +  6KC1  +  3CO2 

(2)  Al(OH),  +  KOH  —  KA10,  +  2H2O 

or  A1C1,  +  4KOH  =  KA1O3  +  3KC1  +  2H2O 

(3)  2KA102  +  H2S  +  2H20  =  2A1(OH)3  +  K2S 

(4)  2KA102  +  2NH4C1  +  H20  =  A120(OH),  +  2KC1  +  2NHS 

(5)  2A1C13  +  3BaC03  +  3H20  =  2A1(OH)8  +  3BaCl2  +  3C02 

*  According  to  Langlota  (A.  Ch.,  1856,  (3),  48,  502)  the  precipitate  with  alkali  carbonates 
always  contains  COi.  He  assigns  the  formula  SCAhOsCCh)  +  SCAlsOs.SHaO). 

t  A  solution  of  barium  hydroxide  may  be  used  to  dissolve  the  Al(OH)s  in  separating  from 
Fe(OH)3  and  Cr(OH)s  ;  especially  valuable  in  detecting  the  presence  of  small  amounts  of  alu' 
minum  when  the  reagents  NaOH  and  KOH  contain  aluminum  (Neumann,  M.,  1894,  15,  53). 


146  ALUMINUM.  §124,  6b. 

ft-  —  Oxalates  do  not  precipitate  aluminum  salts.  The  acetate  of  alum- 
inum is  decomposed  upon  boiling,  forming  the  insoluble  basic  acetate 
(separation  of  iron  and  aluminum  from  the  fourth  group)  :  A1(C2H302)3  + 
H20  =  A1(C2H302)2OH  +  HC2H,02  .  •  The  basic  acetate  is  best  formed  as 
follows:  To  the  solution  of  aluminum  salt  add  a  little  sodium  or  am- 
monium carbonate,  as  much  as  can  be  added  without  leaving  a  precipitate 
on  stirring,  then  add  excess  of  sodium  or  ammonium  acetate,  and  boil  for 
some  time,  when  the  precipitation  at  length  becomes  very  nearly  complete. 

Phenyl  hydrazine,  C6H-NHNH2  ,  completely  precipitates  aluminum  as 
the  hydroxide  from  the  neutral  solution  of  its  salts  (complete  separation 
of  aluminum  and  chromium  from  iron  which  should  be  in  the  ferrous 
condition)  (Hess  and  Campbe.ll,  J.  Am.  Soc.,  1899,  21,  776). 

c.  —  Nitric  acid  is  a  very  poor  solvent  for  metallic  aluminum,  but  a  good 
solvent  for  the  oxide  and  hydroxide.  The  metal  dissolves  in  a  solution  of  the 
normal  aluminum  nitrate,  evolving  hydrogen  and  forming  the  basic  nitrate 
A14O6(N08),  (Ditte,  C.  r.,  1890,  110,  782). 

d.  —  Alkali  phosphates  precipitate  aluminum  phosphate,  A1P04  ,  white, 
insoluble  in  water  and  acetic  acid,  soluble  in  mineral  acids,  and  in  the 
fixed  alkalis  (separation  from  FeP04)   (Grueber,  Z.  angew.,   1896,   741). 
A  separation  of  Al  and  P04  may  be  effected  by  dissolving  in  hydrochloric 
acid  adding  tartaric  acid  and  then  ammonium  hydroxide,  and  digesting 
some  time  with  magnesia  mixture  (magnesium  sulphate  to  which  sufficient 
ammonium  chloride  has  been  added  so  that  no  precipitate  is  obtained 
when  rendered  strongly  alkaline  with  ammonium  hydroxide).     The  filtrate 
contains  nearly  all  of  the  aluminum.     The  same  method  may  be  employed 
with  Fe'"  and*P04  .     See  also  7. 

e.  —  The  sulphide  of  aluminum  cannot  be  prepared  in  the  wet  wa}r,  that 
prepared  in  the  dry  way  being  decomposed  by  water  (Curie,  C.  N.9  1873, 
28,  307).     Hydrosulphuric  acid  does  not  precipitate  aluminum  from  acid 
or  neutral  solutions;  from  its  solutions  in  the  fixed  alkalis  it  is  precipitated 
as  the  hydroxide  on  addition  of  sufficient  hydrosulphuric  acid  to  neutralize 
the  fixed  alkali  (distinction  from  zinc  which  is  rapidly  precipitated  from 
its  alkaline  solutions,  as  the  sulphide).     The  alkali  sulphides  precipitate 
aluminum  from  its  solutions,  as  the  hydroxide;  from  acid  or  neutral  solu- 
tion H2S  is  evolved:   2A1C1,  +  3(NH4)2S  +  6H20  =  2A1(OH)3  +  6NH4C1 
-f-  3H2S  ,  from  solutions  in  the  fixed  alkalis  ammonia  is  evolved,  fixed 
alkali  sulphide  being  formed:  2KA100  +  (NH4)0S  +  2H00  =  2A1(OH)3  + 
KS 


Sodium  thio  sulphate  precipitates,  from  aluminum  salts,  in  neutral  solutions, 
aluminum  hydroxide  with  free  sulphur  and  liberation  of  sulphurous  anhydride: 
2A12(S04)3  +  GNa,S2O3  +  6H2O  =  4A1(OH)3  +  3S2  +  6Na,SO4  +  6SO,  .  A 
small  amount  of  sodium  tetrathionate  is  formed  and  also  some  hydrosulphuric 
acid  (Yortmann,  B.,  1889,  22,  2307).  Sodium  sulphite  also  precipitates  alu- 


§124,  0.  ALUMINUM.  147 

minum  hydroxide,  with  liberation  of  sulphur  dioxide:  2A1CL  +  3Na2SO3  + 
3H2O  =  2A1(OH)3  +  GNaCl  +  ?>SO,  .  Neither  of  1he  above  reagents  precipi- 
tate  iron  salts,  thus  effecting-  a  separation  of  aluminum  (and  chromium)  from 
iron. 

Aluminum,  chromium  and  ferric  sulphates  crvsialli/c  wilh  the  sulphates 
of  the  alkali  metals,  forming  a  class  of  compounds,  AI.TMS,  of  which  th« 
potassium  aluminum  compound  is  perhaps  best  known,  KA1(S04)2.12H20  , 
connnoii  alum.  These  compounds  melt  in  their  water  of  crystallization, 
becoming  anhydrous  upon  further  heating.  The  freshly  ignited  alum  is 
only  sparingly  soluble  in  cold  water,  but  upon  standing  becomes  readily 
soluble,  dissolving  in  less  than  one  part  of  hot  wate1*.  The  alums  are  usu- 
ally less  soluble  than  their  constituent  sulphates  and  may  be  precipitated 
by  adding  a  saturated  solution  of  alkali  sulphate  to  a  very  concentrated  so- 
lution of  Al ,  Cr'"  ,  or  Fe'"  sulphate. 

f. — Aluminum  chloride  is  a  very  powerful  dehydrating-  agent  and  is  much 
used  in  organic  chemistry  as  a  halogen  carrier.  An  impure  aluminum  chlorate, 
mixture  of  KC1O3  and  AL(S04)3  ,  is  much  used  in  calico  printing  (Schlum- 
berger,  Dingl.,  1873,  207,  03).  y. — Aluminum  salts  are  precipitated  by  solu- 
tions of  alkali  arsenites  and  arsenates,  but  not  by  arsenous  or  arsenic  acids. 
7i.— Potassium  chromate  forms  a  yellow  gelatinous  precipitate,  potassium 
bichromate  gives  no  precipitate  with  aluminum  salts,  i. — Solution  of  borax 
precipitates  an  acid  aluminum  borate,  quickly  changed  to  aluminum  hydroxide. 

7.  Ignition. — Compounds  of  aluminum  are  not  reduced  to  the  metal,  but 
most  of  them  are  changed  to  the  oxide,  by  ignition  on  charcoal.  If  now  this 
residue  is  moistened  with  solution  of  cobaltous  nitrate,  and  again  strongly 
ignited,  it  assumes  a  blue  color.  This  test  is  conclusive  only  with  infusible 
compounds,  and  applies  only  in  absence  of  colored  oxides.  Aluminum  com- 
pounds ignited  on  charcoal  in  presence  of  sulphur  are  changed  to  A12S3  (Buch- 
erer,  Z.  angew.,  1892,  483). 

To  separate  Al  from  PO4  ,  fuse  the  precipitate  or  powdered  substance  with 
iy2  parts  finely  divided  silica  and  6  parts  dried  sodium  carbonate  in  a  platinum 
crucible,  for  "half  an  hoiir.  Digest  the  mass  for  some  time  in  water;  add 
ammonium  carbonate  in  excess,  filter  and  wash.  The  residue  consists  of 
aluminum  sodium  silicate;  the  solution  contains  the  PO4  ,  as  sodium  phosphate. 
The  Al  can  be  obtained  from  the  residue  by  dissolving  it  in  hydrochloric  acid, 
evaporating  to  dryness  to  render  the  silica  insoluble.  Treat  with  hydrochloric 
acid  and  filter;  the  filtrate  containing  aluminum  chloride. 

8.  Detection. — After  the  removal  of  the  first  two  groups  it  is  precipi- 
tated with  Cr  and  Fe'"  as  the  hydroxide,  A1(OH)3 ,  by  NH4OH  in  the  pres- 
ence of  NH4C1 .     It  is  separated  from  Fe(OH)3  and  Cr(OH),,  by  boiling 
with  KOH  or  NaOH  or  by  fusion  with  an  alkaline  oxidizing  agent  such  as 
NaCIO  or  Na202 .     From  the  nitrate  acidulated  with  HC1  it  is  precipitated 
as  hydroxide  with    (NH4)2C03  ;  or  it  is   precipitated   from  the   alkaline 
solution  by  an  excess  of  NH4C1  (6a). 

9.  Estimation. — Aluminum   is   usually   weighed   as   the   oxide,   after  ignition 
It  is  separated  from  zinc  as  a  basic  acetate;    from  chromium  by  oxidizing  the 
latter  to  chromic  acid,   by  boiling  with  potassium  chlorate  and  nitric  acid,   or 
by  fusing  with  KNO3  and  Na2CO3  ,    or  by  action  of  Cl  or  Br    in    presence    of 
KOH  ,   or  by  NaaO3  fused  or  in  solution,   and  after  acidulating  with  HC1  precipi- 


CHROMIUM.  §124,  10. 

tating  the  aluminum  with  ammonium  hydroxide.  It  may  be  separated  from  iron 
by  boiling  with  KOH  (6a),  by  Na2S2O3  (6e),  or  by  phenylhydrazine  (6b).  It 
is  separated  from  iron  by  conversion  into  the  oleate  and  dissolving  the  oleate 
of  iron  (Fe'"  or  Fe")  in  petroleum  (Borntraeger,  Z.,  1893,  32,  187).  It  is  some- 
times precipitated  and  weighed  as  the  phosphate. 

10.  Oxidation. — Aluminum  reduces  solutions  of  Pb  ,  Ag  ,  Hg  *,  Sn  ,  Bi 
(incompletely),  Cu  f,  Cd  ,  Co ,  Ni ,  Zn  J  and  Gl  (in  alkaline  mixture  only), 
Te,  Se,  An,  and  Pt ,  to  the  metallic  state;  ferric  salts  to  ferrous  salts; 
As  and  Sb  with  HC1  become  respectively  AsH3  and  SbH3  ,with  alkalis  As'" 
is  reduced  to  AsH3 .  Asv  is  unchanged  (§69,  6'&  and  10),  and  Sb"'  and 
Sbv  become  Sb°.  Aluminum  salts  are  not  reduced  to  the  metallic  state 
by  any  other  compounds  at  ordinary  temperature;  by  fusion  with  K  or  Na 
metallic  aluminum  is  obtained,  much  better,  however,  by  the  aid  of  the 
electric  current. 

§125.  Chromium.     Cr  =  52.0.     Valence  two,  three  and  six. 

1.  Properties. — Specific    gravity,    6.92    (Moissan,    C.    r.,    116).     Melting  point 
1520°  (Cir.  B.  £.,  36,  1915).     A  grayish-white  crystalline  metal.     The  hardness 
of  steel  is  greatly  increased  by  the  presence  of  less  than  one  per  cent  of  chromium. 
It  is  non-magnetic  (Woehler,  A.,  1859,  111,  231).     It  burns  to  the  oxide  Cr2O3 
when  heated  to  200°  to  300°  in  the  air  (Moissan,  C.  r.,  1879,  88,  180). 

2.  Occurrence. — Not  found  native.     Chrome-ironstone  or  chromite  (FeOCr2O3) 
is  the  chief  ore  of  chromium,  and  is  usually  employed  in  the  manufacture  of  chro- 
mium  compounds.     Chromite   and   also   Daubreelite    (FeCr2S4)  ,    are   frequently 
found  in  meteorites;    it  also  occurs  in  crocoite   (PbCrO4)  ,  and  other  rare  chro- 
mates  and  sulphates. 

3.  Preparation. —  (1)    By    electrolysis    of    the    chloride.     (2)    By    fusing    the 
chloride  with  potassium  or  sodium.     (3)  By  ignition  of  the  oxide  with  carbon. 
(4)  By  fusing  CrCl3  with  Zn  ,  Cd  or  Mg  ,  using  KC1  and  NaCl  as  a  flux,  and 
removing  the  excess  of  the  Zn  ,  Cd  or  Mg  by  dissolving  in  nitric  acid,  which 
does  not  dissolve  metallic  chromium.     (J)   By  ignition  of  the  oxide  with  alu- 
minum (Goldschmidt,  A.,  1898,  301,  19). 

4.  Oxides  and  Hydroxides. — CJiromous  oxide,  CrO  ,  has  not  been  isolated.     The 
corresponding   hydroxide,    Cr(OH)2  ,    is    made    by    treating    CrCL    with    KOH. 
Chromic  oxide,  Cr2O3  ,  is  made  by  a  great  variety  of  methods,  among  which  are 
fusing  the  nitrate,  or  higher  or  lowrer  oxides  and  hydroxides  in  the  air;  heating 
mercurous  chromate,  or  the  dichromates  of  the  alkalis: 

4Hg2Cr04  =  2Cr203  +  SHg  +  502 
(NH4)2Cr207  =  Cr203  +  N2  +  4H20 
4K2Cr2O7  =  2Cr,O3  +  4K2Cr04  +  3O2 

In  the  last  the  K2CrO4  may  be  separated  by  water.  After  heating  to  redness, 
Cr,O3  is  insoluble  in  acids.  Chromic  hydroxide,  Cr(OH)3  ,  is  precipitated  by 
adding  NH4OH  to  chromic  solutions.  That  formed  by  precipitating  with  KOH 
or  NaOH  retains  traces  of  the  alkali,  not  easily  removed  by  washing. 

Chromium  trioxide  or  chromic  anhydride,  Cr03  ,  is  formed  as  brown-red 
needles  upon  addition  of  concentrated  sulphuric  acid  to  a  concentrated  solution 
of  K2Cr2O7;  to  be  freed  from  sulphuric  acid  it  must  be  recrystallized  from 
water,  in  which  it  is  readily  soluble,  or  treated  with  the  necessary  amount  of 

*  Klandy,  C.  C.,  1893,  201 ;  Wislicenus,  B.  1895,  28, 1323.  t  Tommasi,  Bl.,  1882,  (2),  37.  .{,•:}. 

t  Flavitsky,  B.,  1873,  6, 195 ;  Zimmerman,  Z.,  1888,  27,  61? 


,  60.  CHROMIUM.  149 

BaCrO4  (Moissan,  A.  Ch.,  1885,  (6),  6,  568).  It  is  also  prepared  by  transposi- 
tion of  BaCrO4  with  HNO3  or  H2SO, ;  PbCrO4  with  H2SO4  ;  and  Ag2CrO4  with 
HC1;  etc.  It  melts  at  196°,  decomposing  at  higher  temperature  into  Cr2O3 
and  O  .  It  is  used  in  dyeing  silk  and  wool,  but  not  cotton  fabrics.  It  is  a 
powerful  oxidizing  agent,  being  reduced  to  chromic  oxide.  The  existence  of 
chromic  acid,  H2CrO4 ,  is  disputed  (Moissan,  L  c.;  Field,  C.  N.,  1892,  65,  153; 
and  Ostwald,  Zeit.  phys.  Ch.,  1888,  2,  78).  Two  series  of  salts  are  formed  as  if 
derived  from  chromic  acid,  H2CrO4  ,  and  dichromic  acid,  H2Cr2O7  .  The  salts 
are  quite  stable  and  find  an  extended  application  in  analytical  chemistry  (Qh. 
§57,  §59,  §186,  etc.). 

5.  Solubilities. — a. — Metal. — Chromium  is  not  at  all  oxidized  by  water  or 
moist  air  at  100°.  Heated  above  200°  it  is  oxidized  to  Cr,O,  ,  rapidly  in  pres- 
ence of  KOH  .  It  is  soluble  in  HC1  or  dilute  H,S04;  insoluble  in  concentrated 
HoSO4  or  in  HNO;i  ,  dilute  or  concentrated.  Chlorine  or  bromine  attack  it 
with  formation  of  the  corresponding-  halides  (Woehler,  L  c.;  Ufer,  A.,  1859,  112, 
302).  &. — Oxides  and  Hydroxides. — Chromic  o.ride,  Cr,03  ,  is  insoluble  in  water, 
slowly  soluble  in  acids,  but  not  at  all  if  previously  ignited  (Traube,  A.,  1848, 
66,  88);  the  hydroxide  is  insoluble  in  water,  soluble  in  acids,  sparingly  soluble 
in  ammonium  hydroxide,  soluble  in  fixed  alkalis  to  chromites,  reprecipitated 
again  upon  boiling-.  The  presence  of  other  metallic  hydroxides,  as  iron,  etc., 
hinders  the  solution  in  fixed  alkalis.  Chromic  anhydride,  Cr03  ,  is  very  soluble 
in  water,  soluble  in  reducing  acids  to  chromic  salts. 

c. — Salts. — Chromic  sulphide  is  not  formed  in  the  wet  way,  being 
decomposed  by  water;  the  phosphate  is  insoluble  in  water.  The  chloride 
exists  in  two  modifications ;  a  deliquescent  soluble  chloride,  which  also 
forms  a  soluble  basic  chloride  (Ordway,  Am.  S.,  1858  (2),  26,  202); 
and  a  violet  sublimed  chromic  chloride  absolutely  insoluble  in  water, 
hot  or  cold,  or  in  dilute  or  concentrated  acids,  the  presence  of  a  very 
small  amount  of  chromous  or  stannous  chloride  at  once  renders  this  modi- 
fication soluble  in  water  (Peligot,  A.  Ch.,  1846  (3),  16,  298);  the  bromide 
and  sulphate  also  exist  in  soluble  and  insoluble  modifications;  the  nitrate 
and  also  the  basic  nitrates  are  readily  soluble  in  water  (Ordway,  1.  c.). 
There  are  many  double  salts,  the  sulphates  of  chromium  and  the  alkali 
metals,  chrome  alum,  forming  salts  similar  to  the  corresponding  aluminum 
compounds.  There  are  two  modifications  of  solutions  of  chromium  salts, 
one  having  a  green  color  and  the  other  violet  to  red,  the  tints  being 
modified  somewhat  by  the  degree  of  the  concentration.  All  normal  chromic 
salts  in  solution  have  an  acid  reaction,  being  partially  hydrolized. 

6.  Reactions.*  a. — Alkali  hydroxides  and  carbonates  precipitate  solu- 
tions of  chromic  salts,  as  chromium  hydroxide,  gelatinous,  gray-green  or 
gray-blue  according  to  the  variety  of  solution  from  which  it  is  obtained 
(5c),  insoluble  in  water,  soluble  in  acids;  soluble  in  excess  of  the  fixed 
alkalis  to  chromites :  Cr(OH)3  +  KOH  <=>  KCr02  +  2H20  .  This  reaction 

*  Chromous  salts  are  very  unstable,  they  are  great  reducing  agents,  oxidizing-  rapidly  when 
exposed  to  the  air.  They  are  almost  n°ver  met  with  in  analysis.  Chromous  chloride,  CrCl2,  is 
formed  when  the  metal  is  heated  in  contact  with  hydrochloric  acid  gas  (Ufer,  I.  c  );  also  by  re- 
duction of  CrCI3  with  hydrogen  in  a  heated  tube  (Moberg,  J.  pr.,  1848,  44r,  322).  Precipitates  are 
formed  in  its  solutions  by  the  alkali  hydroxides,  carbonates,  sulphides,  etc.  (Moissan,  J37.,  1HX2 
(2),  37,  296). 


150  CHROMIUM.  §125,  6£. 

is  reversible.  The  KOH  tends  to  form,  while  the  water,  especially  when 
hot,  tends  to  decompose  the  chromite.  The  chromium  may  therefore  be 
reprecipitated  from  this  solution  if  the  excess  of  KOH  is  small  and  the 
solution  is  diluted  and  boiled.  (Distinction  from  aluminum.)  As  ammo- 
nium chloride  reacts  with  the  caustic  potash  forming  potassium  chloride 
and  the  weak  alkali,  ammonium  .hydroxide,  the  chromium  may  also  be  pre- 
cipitated by  the  addition  of  ammonium  chloride  and  heating.  The  presence 
of  ferric  hydroxide  and  some  other  compounds  greatly  hinders  the  solution 
in  fixed  alkalis,  hence  chromium  cannot  be  separated  from  iron  by  excess 
of  fixed  alkali.  Chromium  hydroxide  is  slightly  soluble  in  excess  of  cold 
ammonium  hydroxide  to  a  violet  solution,  completely  reprecipitated  on 
boiling.  The  precipitate  formed  with  the  alkali  carbonates  is  almost 
entirely  free  from  carbonate:  2CrCl3  +  3Na,C03  +  3H20  =  2Cr(OH)3  + 
GNaCl  -(-  3C02 .  Barium  carbonate  precipitates  chromium  from  its  solu- 
tions (better  from  the  chloride)  as  a  hydroxide  with  some  basic  salt,  the 
precipitation  being  complete  after  long  digestion  in  the  cold  (separation 
from  the  fourth  group).  For  removal  of  excess  of  reagent,  add  H0S04 
and  the  filtrate  will  contain  the  chromium  as  a  sulphate. 

Alkali  dichromates  are  changed  to  normal  chromates  by  alkali  hydrox- 
ides or  carbonates. 

1). — Chromium  forms  no  basic  acetate  and  remains  in  solution  when  the 
basic  acetates  of  aluminum  and  ferric  iron  are  formed  (6&,  §124  and  §126). 
Potassium  cyanide  precipitates  chromium  hydroxide.  Oxalates  and  ferro- 
cyanides  cause  no  precipitate.  H2Cr04  is  reduced  to  chromic  compounds 
by  K3Fe(CN)6  and  KCNS.  r. — Nitrites  or  nitrates  are  without  action  upon 
chromium  salts  in  the  wet  way,  but  upon  fusion  in  presence  of  nitrites  or 
nitrates  and  alkali  carbonate  a  chromate  is  formed  (separation  from  Fe  and 
Al).  d. — Hypophosphorous  acid  reduces  chromates  to  chromic  salts.  Soluble 
phosphates,  as  Na,HPO4  ,  precipitate  chromic  phosphate,  CrPO4  ,  insoluble  in 
acetic  acid,  decomposed  by  boiling1  with  KOH  ,  leaving  the  phosphate  in  solu- 
tion (Kammerer,  J.  C.,  1874,  27,  1005). 

e. — Hydrosiilplmric  acid  is  without  action  upon  neutral  or  -acid  solutions 
of  chromic  salts,  chromites  as  KCr02  are  precipitated  as  chromium 
hydroxide;  2KCr02  +  HJ3  +  2H20  ="  2Cr(OH)3  +  K2S .  The  hexad 
chromium  of  chromates  is  reduced  to  the  triad  condition  with  liberation 
of  sulphur,  in  neutral  or  alkaline  solutions,  chromium  hydroxide  being 
formed:  2K2Cr,0-  +  8H2S  =:  4Cr(OH)3  +2K2S  +  3S2  +  2H20  ;  in  acid 
solutions  a  chromic  salt  is  formed  (10).  Alkali  sulphides  precipitate 
chromic  salts  as  the  hydroxide  liberating  H2S  : 

2CrCL,  +  3(NH4)»S  +  GH20  =  2Cr(OH)3  +  GNH4C1  +  ?>H2S 
Chromates  are  reduced  and  precipitated  as  chromium  hydroxide  with  sepa- 
ration of  sulphur:    4K2Cr04  +  6(NH4)2S  +  4H20  =  4Cr(OH)3  +  8KOH 


§125,  8.  CHROMIUM. 

+  3S2  -f-  12NH3  .  Soluble  sulphites  and  thiosulpliates  reduce  chromates 
in  acid  solution  (Donath,  /.  C.,  1879,  36,  401;  Longi,  Gazzetta,  1896,  26, 
ii,  119). 

f. — Hydrochloric  acid  reduces  chromates  to  chromic  chloride  on  boiling,, 
with  evolution  of  chlorine:  2K,Cr04  +  1GHC1  =  2CrCl3  -f  4KC1  +  3C12  + 
SHoO  ;  more  readily  without  evolution  of  chlorine  in  presence  of  other 
easily  oxidized  agents,  as  alcohol,  oxalic  acid,  etc. :  K2Cr207  -f-  8HC1  -f- 
3C2H8OH  =  2KC1  -f  2CrCl3  +  3C2H40  (acetaldehyde)"+"?'H20  .  If  the 
dry  chromate  be  heated  with  sulphuric  acid  and  a  chloride  (transposable 
by  sulphuric  acid)  (§269,  5),  brown  fumes  of  chromium  dioxydichloride 
are  evolved:  K2Cr207  -f  4NaCl  -f-  3H2S04  *=  2Cr02Cl2  +  K2S04  +  2Na2S04 
+  3H20  (§269,  Sd)  (Moissan,  Bl,  1885  (2),  43,  6).  To  obtain  a  quantity  of 
Cr02Cl2 ,  Thorpe  (J.  C.,  1868,  21,  514)  recommends  10  parts  of  NaCl  and 
12  parts  K2Cr,07  fused  together  and  distilled  with  30  parts  of  H2S04 . 
Hydrobromic  acid  reduces  chromates  to  chromic  bromide  with  evolution 
of  bromine;  hydriodic  acid  to  chromic  iodide  with  evolution  of  iodine. 
In  the  presence  of  hydrochloric  or  sulphuric  acids  all  the  bromine  or 
iodine  is  set  free.  K2Cr207  +  6HI  +  4H2S04  =  K2S04  +  Cr2(S04)3  4- 
3I2  +  7H20  .  Hydriodic  acid  acts  most  readily  upon  chromates,  the 
hydrochloric  least  readily.  Chromic  hydroxide  and  chromic  salts,  when 
boiled  with  chloric  or  bromic  acids,  or  potassiuni^hlorate  or  brornate  and 
nitric,  sulphuric  or  phosphoric  acids,  become  chromic  acid. 

g. — Soluble  arsenites  and  arsenates  form  corresponding'  salts  with  chromic 
salts.  Chromates  in  acid  solution  are  instantly  reduced  to  chromic  salts  by 
arsenites  or  arsenous  acid.  Chromic  acid  boiled  with  arsenous  acid  in  excess 
gives  CrAsO,  (Neville,  J.  C.,  1877,  31,  283). 

k. — Potassium  chromate  colors  an  acid  solution  of  chromic  salt  brown-yellow' 
on  addition  of  ammonium  hydroxide,  a  precipitate  of  the  same  color  is  obtained, 
chromic  chromate  (Maus,  Po#</.,  1827,  9,  127).  The  alkali  metals  form  two 
classes  of  chromates:  yellow  normal  chromates  and  reddish  dichromates 
(Schulernd,  •/.  C.,  1879,  36,  298).  The  chromates  of  the  alkalis,  and  those  of 
magnesium,  calcium,  zinc  and  copper  are  soluble;  those  of  strontium,  mercury 
(Hg")  are  sparingly  soluble;  and  those  of  barium,  manganese,  bismuth,  mer- 
cury (Hg'),  silver  and  lead  are  insoluble  in  water.  Alkali  chromates  or 
dichromates  are  precipitated  as  normal  chromates  (in  some  cases  as  dichro- 
mates) (Preis  and  Kayman,  B.,  1880,  13,  340)  by  solutions  of  silver,  lead,  mer- 
cury (HgO  and  barium  salts.  Silver  chromate  is  dark  red,  soluble  in  nitric 
acid  and  ammonium  hydroxide  (§59,  67?) ;  lead  chromatc  is  yellow,  transposed 
with  difficulty  by  nitric  acid  (Duvillier,  A.  C7i,,  1873,*  (4),  30,  212),  insoluble  in 
acetic  acid  (§57,'  67/):  barium  chromate,  yellow,  is  soluble  in  hydrochloric  and 
nitric  acids,  sparingly  soluble  in  chromic  acid  (§186,  67»-). 

7.  Ignition.— Chromic  oxide,  chromic  salts  and  chromates  dissolve  in  beads 
of  microcosmic  salt,  and  of  borax,  before  the  blow-pipe,  in  both  reducing  and 
oxidizing  flames,  with  a  yellowish-green  tint  while  hot,  becoming  emerald 
green  when  cold.  By  ignition  on  charcoal  the  carbon  deoxidizes  chromic 
anhydride,  CrO3  ,  free  or  combined,  and  a  green  mass,  Cr2O3  ,  is  left.  When 
chromium  compounds  are  fused  with  an  alkali  carbonate,  and  a  nitrite,  nitrate^ 
chlorate,  bromate  or  iodate,  an  alkali  chromate  is  formed,  soluble  in  water 
(distinction  from  Al  and  Fe). 

8.  Detection, — If  present  as  chromate  (solution  red  or  yellow),  it  is 


152  CHROMIUM.  §125,  9. 

reduced  by  HC1  and  alcohol  or  by  H2S.  Precipitated  with  Fe"'  and  Al  , 
after  the  removal  of  the  metals  of  the  first  and  second  groups,  by  NH4OH 
in  presence  of  NH4C1  .  Boiling  with  KOH  separates  the  Al  and  leaves  the 
Cr  and  the  Fe  ,  as  hydroxides.  The  precipitate  is  fused  on  a  platinum  foil 
with  Na2C03  and  KN03  which  oxidizes  the  Cr  to  an  alkali  chromate,  soluble 
in  water  (separation  from  the  Fe).  The  Cr  is  identified  after  acidulation 
with  HCoH302  by  the  formation  of  the  yellow  lead  chromate,  using 
Pb(C.>H302)2  .  These  metals  may  also  be  separated  by  the  methods  given 
in  10". 

9.  Estimation.  —  Chromium  is  usually  estimated  gravimetrically    (1)   as  the 
oxide.     It  is  brought  into  this  form  either  by  precipitation  as  a  hydroxide  (6a) 
and  ignition  or,  in  many  cases,  by  simple  ignition  (4).     (2)  As  chromate,  it  may 
be  precipitated  with  barium  chloride,  dried  and  weighed  as  such;  or  in  acetic 
acid   solution  it  may  be  precipitated   as   PbCrO4    by  PbCC.HgO,),  ,   dried   and 
weighed.     Volumetrically,  as  a  chromate  (if  present  as  chromic  salt  it  may  be 
oxidized  to  a  chromate).     (3)  By  titration  with  a  standard  solution  of  ferrous 
sulphate.     (//)  By  liberation  of  iodine  from  hydriodic  acid  (60)  and  measuring 
the  amount  of  iodine  liberated  with  standard  sodium  thiosulphate  solution. 

10.  Oxidation.  —  Chromous  compounds  are  very  strong  reducing  agents, 
changing  HgCl2  to  HgCl  ,  CuS04  to  Cu°,  SnCl2  to  Sn°,  etc.     Chromic  com- 
pounds are  oxidized  to  chromates  by  chlorates  (Giacomelli,  UOrosi,  1895, 
18,  48;  Storer,  Am.  S.,  1869,98,190)  (6/),  Na,02,  Mn02  (Marchal  and  Wier- 
nick,  Z.  angew.,  1891,  511),  and  Pb02  in  acid  solution;  in  alkaline  mixture, 
by  reducing  Pb02  to  PbO  ,  Ag20  to  Ag°,  Hg20  and  HgO  to  Hg°,  CuO  to 
CiuO  ,  KMn04  and  K2Mn04  to  Mn02  (Donath  and  Jeller,  C.  C.,  1887,  151); 
by  Cl,  Br,  and  I,  forming  the  corresponding  halide;   and  by  H202* 

(Baumann,  Z.  angew.,  1891,  139). 

The  halogens  in  alkaline  solution  may  be  used  to  separate  chromium  from 
iron  and  aluminum.  The  halogens  react  with  the  alkali  and  oxidize  the 
chromium  to  chromate  according  with  the  following  reactions: 

2NaOH  +  C12  =  NaCl  +  NaCIO  +  H2O 

2NaCrO2  +  SNaCIO  +  NaOH  =  2Na2CrO4  +  3NaCl  +  H2O. 


The  ferric  hydroxide  is  not  acted  upon,  while  the  aluminum  hydroxide 
dissolves  in  the  fixed  alkali  and  may  be  separated  from  the  chromium  by 
the  addition  of  ammonium  chloride  and  warming.  Na202  or  H202  in  alka- 
line solution  produce  a  similar  reaction. 

2NaCrO2  +  SNa^  +  2H2O  =  2Na2CrO4  +  4NaOH, 

the  separation  of  iron  and  aluminum  being  effected  by  the  method  already 
given. 

*  The  use  of  H2O2  in  alkaline  solution  is  proposed  by  Riggs  (Am.  £,  1894,  148,  409)  in  the  sepa- 
ration of  Al,  F'e  and  Cr.  100  cc.  water,  10  cc.  H2O2,  and  one  gram  of  BfaOH  are  added  to  the 
freshly  precipitated  hydroxides  and  digested  until  effervescence  ceases.  Filter  off  the  precipi- 
tate of  ferric  hydroxide,  acidify  tho  filtrate  with  acetic  acid  and  precipitate  the  aluminum  witb 
ammonium  hydroxide.  Tfce  chromium  if  present  will  be  in  the  filtrate  as  sodium  ohromatc. 


§126,  1.  IRON.  153 

A  chromate  is  also  formed  when  chromium  compounds  are  fused  with 
an  alkali  carbonate  and  an  oxidizing  agent  (7). 

For  this  purpose  sodium  carbonate  and  potassium  nitrate  or  sodium 
peroxide  are  frequently  used.  The  following  reactions  take  place: 

2Cr(OH)3  +  2Na2CO3  +  3KNO3  =  2Na2CrO4  +  3KNO2  -f  2CO2  +  3H2O 

2Cr(OH)3  +  3Na2O2  =  2Na2CrO4  +  2NaOH  +  2H2O. 

Ferric  hydroxide  is  not  acted  upon  while  aluminium  hydroxide  is  con- 
verted into  aluminate. 

2A1(OH)3  +  Na2CO3  =  2NaAlO2  +  CO2  +  3H2O  . 

On  dissolving  the  fused  mass  in  water,  the  three  metals  may  be  separated 
by  the  method  already  given. 

Chromic  oxide  (not  ignited)  or  chromic  chloride  at  440°  in 
a  current  of  chlorine  become  Cr02Cl2  (Moissan,  El.,  1880,  (2),  34,  70). 
Chromic  acid  and  chromates  are  reduced  to  chromic  compounds  by 
H2C204  (Werner,  /.  C.,  1888,  53,  602),  K4Fe(CN)6 ,  KCNS ,  H2S ,  (NH4)2S , 
Na2S20, ,  S02 ,  H202 ,  etc.  Of  most  common  occurrence  in  qualitative 
analysis  is  the  action  of  hydrosulphuric  acid  and  alkali  sulphides;  at  first 
sulphur  is  liberated,  a  part  of  which  may  be  oxidized  to  sulphurous  and 
sulphuric  acids  (Parsons,  C.  N.,  1878,  38,  228). 

2K2Cr207  +  16HC1  +  6H2S  =  4CrCl3  +  4KC1  +  3S2  +  14H2O 
12H2Cr04  +  3S2  =  4Cr20,Cr04  +  6SO,  -f  12H2O 
2H2Cr04  +  3SO,  =  Cr2(S04)3  +  2H20 

While  H202  in  alkaline  solution  oxidizes  Cr'"  to  Crvl,  in  acid  solution  the 
reverse  action  takes  place:  2H2Cr04  +  3H2S04  -f-  3H202  =  Cr2(S04)3  + 
:>02  +  8H20  (Baumann,  1.  c.).  With  chromate  in  acid  solution,  the  H,02 
at  first  gives  a  deep  blue  solution  (probably  of  the  very  unstable  perchromic 
acid,  HCr04),  followed  by  the  reduction  to  a  chromic  salt.  The  blue  color 
gives  a  very  delicate  test  for  chromium.  The  test  is  rendered  more  deli- 
cate by  the  addition  of  a  few  c.c.  of  ether  and  shaking.  The  ether  dis- 
solves the  blue  compound  and  forms  a  blue  layer  on  standing.  One  part 
of  chromic  acid  in  40,000  parts  of  water  can  be  detected  by  this  reaction. 
Vanadic  acid  interferes  with  the  delicacy  (Reichard,  Z.,  40,  577). 

§126.  Iron,  (Ferrum).    Fe  —  55.84.    Usual  valence  two  and  three. 

1.  Properties. — Specific  gravity,  variable,  depending  upon  the  purity  and 
methods  of  preparation.  7.85  at  16°  (Caron,  C.  r.,  1870,  70,  1263),  8.139 
(Chandler-Roberts,  C.  N.,  1875,  31,  137).  Melting  point,  cast  iron,  1130°  to 
1375°;  steel,  1375  to  <1530°;  pure  iron,  1530°  (Cir.  B.  S.,  36,  1915).  Pure  iron 
is  silver-white,  capable  of  taking  a  remarkably  fine  polish;  it  is  among  the  most 
ductile  of  metals,  in  this  property  being  approached  by  nickel  and  cobalt  (§73,  1); 
it  is  the  hardest  of  the  ductile  metals  (Calvert  and  Johnson,  DingL,  1859,  162, 
129),  and  in  tenacity  it  is  surpassed  only  by  cobalt  and  nickel  (§132,  1).  It 
softens  at  a  red  heat  and  may  be  welded  at  a  white  heat.  Finely  divided  iron 
burns  in  the  air  when  ignited;  that  made  by  reduction  in  hydrogen  may  ignite 
spontaneously  when  exposed  to  the  air.  When  pure  iron  is  heated,  or §  cooled 
through  certain  ranges  of  temperature,  polymorphic  changes  occur,  which  are 


154  IRON.  §126,  2. 

accompanied  by  an  absorption  (on  heating)  or  an  evolution  (on  cooling)  of  heat, 
and  changes  in  the  physical  properties  of  the  iron;  these  polymorphic  modifica- 
tions are  called  Alpha,  Beta  and  Gamma  iron,  respectively. 

Alpha  iron  is  stable  in  all  ranges  of  temperature  up  to  768°  C.  (Ac2)  ;  *  at 
768°  C.  the  change  Alpha  ^  Beta  occurs;  Beta  iron  is  stable  between  768°  C. 
and  909°  C.  (Ac3)  ;  *  at  909°  C.  the  change  Beta  -»  Gamma  takes  place 
(Burgess  and  Crowe,  Reprint,  213,  Bui.  B.  S.,  10,  1913).  Gamma  iron  is  capable 
of  dissolving,  and  retaining  in  solid  solution,  carbon  as  iron  carbide  (Fe?C)  ; 
the  presence  of  Fe3C  in  solid  solution,  however,  progressively  lowers  the  tem- 
perature of  the  A3  *  point,  at  which  the  change  Gamma  — >  Be t-i  occurs,  and 
causes  the  appearance  of  a  third  point  (Ai)  *  at  ±  700°  C.,  the  temperature 
of  which  is  unaffected  by  varying  percentages  of  carbon.  Because  of  the  pro- 
gressive lowering  of  the  temperature  at  which  the  Gamma  — >  Beta  (Ar3)  *  change 
takes  place  by  increasing  percentages  of  carbon,  the  (A3)  *  point  merges  into  the 
(A2)  *  point  at  ±768°  C.,  and  ±0.33%  C.  At  this  temperature  and  concentration, 
and  between  it  and  ±700°  C.  (Ai)*  Gamma  iron  probably  passes  directly  into  the 
alpha  modification,  with  a  continued  concentration  of  carbon  in  the  solid  solution 
still  remaining,  until  the  Ar3-2  *  point  and  ATI  *  point  coincide  (Ar3-2-i).  * 

This  point  (±700°  C.,  and  ±0.835%  C.)  marks  the  lowest  temperature  at  which 
carbon  (as  Fe3C)  will  remain  in  solid  solution;  if  the  temperature  falls  below  Ari,* 
Carbon  separates  from  the  solid  solution  as  Fe3C  (cementite)  and  at  the  same 
instant  the  Gamma  iron,  which  up  to  that  moment  had  acted  as  a  solvent  for  the 
carbon  is  transformed  to  Alpha  iron,  simultaneously  with  the  precipitation  of 
the  cementite,  forming  that  interstratified  conglomerate  of  Alpha  ferrite,  and 
cementite,  known  as  Pearlite  (the  eutectoid). 

Above  ±  700°  C.  the  solubility  of  Fe3C  in  Gamma  iron  increases  with  the  tem- 
perature, reaching  a  maximum  at  a  concentration  of  about  1.7%  carbon  and  a 
temperature  of  1130°  C.  Hence  all  slowly  cooled  iron-carbon  alloys  containing 
less  than  the  eutectoid  percentage  of  carbon  (<0.835%  C.;  hypo-eutectoid 
steels)  will  consist  of  structurally  free  Alpha  Ferrite  (pure  iron)  and  Pearlite; 
those  with  more  than  the  eutectoid  percentage  of  carbon  (between  0.835%  C. 
and  1.7%  C.;  hyper-eutectoid  steels)  will  consist  of  structurally  free  Cementite 
(Fe3C)  and  Pearlite;  those  of  eutectoid  composition  (0.835%  C.;  eutectoid 
steel),  will  consist  of  Pearlite  only.  When  such  steels,  provided  the  carbon  con- 
tent lies  between  ±0.25%  and  1.75%,  are  heated  to  a  temperature  higher  than 
Ac3  *  (Accm  *  in  the  case  of  hyper-eutectoid  steels)  and  then  suddenly  cooled 
(e.g.,  by  quenching  in  water),  the  changes  which  would  normally  occur  on  slow 
cooling  through  Ar3,  Ar2,  Ari  *  are  partially  suppressed;  the  steel  becomes  hard 
and  brittle,  the  carbon  being  retained  in  solid  solution.  This  process  is  called 
"hardening."  The  hardened  steel  may  be  "  tempered"  by  reheating  to  temperatures 
lower  than  Aci*  (±700°  C.)  which  causes  certain  structural  changes  to  take  place 
leading  to  the  formation  of  a  series  of  transition  products,  the  general  effect  of  which 
is  to  soften  the  steel  somewhat  and  lessen  its  brittleness  to  the  extent  desired. 

2.  Occurrence. — Iron  is  rarely  found  native  except  in    meteorites;     the    iron 
minerals  used  as  ores  are  hematite  (Fe2Oj),    limonite   (2Fe(OH)3.Fe2Os)  ,-    mag- 
netite (Fe3O4)  ;    to  a  less  extent,  siderite  (FeCO3),    clay  iron  stone  (FeCO.i  with 
clay)  black  band  (FeCO3  with  bituminous  matter);  it  also  occurs  as  pyrite  (FeS2), 
marcasite  (FeS2),  pyrrhotite  (FenSn+i)  and  widely  distributed  in  many  other  min- 
erals and  rocks. 

3.  Preparation. — Pure  iron  is  not  usually    found  in    the    market,   although 
some  of  the  commercial  products  approach  it  very  closely. 

Pig  iron  is  produced  by  smelting  iron  ore  mixed  with  coke  and  limestone  in 
blast  furnaces;  the  resulting  product  is  subjected  to  remelting  processes  (prod- 
uct, gray  and  malleable  iron  castings),  or  to  conversion  processes  (products, 
Bessemer  steel;  open- hearth  steel,  acid  and  basic;  wrought  iron);  pure  iron  may 
be  made  by  electrolysis,  and  by  heating  its  purified  salts  with  hydrogen. 

4.  Oxides  and  Hydroxides. — Ferrous  oxide,  FeO ,  is  made  from  Fe2O3  by  heat- 
ing it  to  300°  in  an   atmosphere  of  hydrogen;    also  by  heating  FeC2O4  to  160°, 

*  Ari-2-2  is  derived  from  Arret  (a  halt,  or  pause)  and  refroidissement  (cooling) ;  hence,  a 
halt  in  cooling  at  certain  critical  temperatures.  Ar3  =898°C.;  Ar2  =768°C.;  An  =  ±700°C. 
Aci-:-3  (c  =  chauffant,  heating)  refers  to  corresponding  points  during  heating.  Acs  is  slightly 
higher  (909°  C.)  than  Ar».  Accra  or  Arem  refers  to  temperatures  at  which  Cementite  dissolves 
in,  or  is  precipitated  from  Gamma  iron  during  heating  or  cooling. 


§126,  51.  IRQ!!.  155 

air  being  excluded.  It  takes  fire  spontaneously  in  the  air,  oxidizing-  to  Fe2O3  . 
Ferrous  li!/<lrt>.ri<I<>,  Fe(OH),  ,  is  formed  by  precipitating-  ferrous  suits  with  KOH 
or  NaOH  ,  perfectly  white  when  pure,  but  usually  green  from  partial  oxidation. 
/•V/r/V  o.r'nle,  Fe,O;i  ,  is  formed  by  heating-  FeO  ,  Fe(OH)2  ,  or  any  ferrous  salt 
consisting-  of  a  volatile  or  organie  acid  in  the  air;  more  rapidly  by  heating 
Fe(OH)3  ,  Fe(NO8),  ,  or  Fe2(S04)3  .  Ferric  hydroxide  is  formed  by  precipitat- 
ing- cold  dilute  ferric  salts  with  alkalis  or  alkali  carbonates,  and  drying  at  100°. 
If  KOH  or  NaOH  is  used,  the  precipitate  requires  longer  washing  than  when 
NH,OH  is  employed.  By  increasing  the  temperature  and  concentration  of  the 
solutions,  the  following  definite  compounds  may  be  formed:  FeO(OH)  , 
Fe,0(OH)4  ,  Fe40,(OH)L,,  Fe4O3(OH),,,  Fe3O2(OH)5  .  Fe3O4  is  slowly  formed 
by  heating-  FeO  or  Fe,O3  to  a  white  heat.  Its  corresponding  hydroxide  may  be 
made  by  precipitation:  FeCL  -f  2FeCl3  +  8NH4OH  =  Fe3(OH)8  +  8NH4C1  . 
Fe3(OH)s  when  heated  to  90°  forms  Fe304  .  The  black  color  and  magnetic 
properties  show  that  it  is  a  chemical  salt  and  not  a  mechanical  mixture  of  FeO 
and  Fe,O3  .  Fe'"  acts  as  an  acid  towards  the  Fe";  this  oxide,  Fe3O4  ,  or 
FeFe_,O4  ,  may  be  called  ferrous  ferrite.  Other  ferrites  have  been  formed,  e.  g., 
calcium  ferrite,  CaFe,O4;  MgFe2O4  and  BaFe2O4  (List,  #.,  1878,  11,  1512);  zinc 
ferrite,  ZnFe,O4  .  Compare  potassium  aluminate,  KA10,  (§124,  6a),  and  potas- 
sium chromite,  KCrO,  (§125,  6a).  Ferric  acid,  HL,Fe64  ,  and  its  anhydride, 
FeO3  ,  have  not  been  isolated.  Potassium  ferrate,  K2Fe04  ,  is  made  (/)  by  elec- 
trolysis; (2)  by  heating  iron-filings,  FeO  or  Fe2O3  ,"  to  a  red  heat  with  KN03; 
(3)  by  heating  Fe(OH)3  with  potassium  peroxide  K202;  (4)  by  passing  Cl  or  Br 
into  a  solution  of  5  parts  of  KOH  in  8  parts  of  water  in  which  Fe(OH)3  is 
suspended;  the  temperature  should  be  not  above  50°.  It  has  a  purple  color;  is 
a  strong  oxidizing  agent.  It  slowly  decomposes  on  standing:  4K,FeO4  + 
10H2O  =  8KOH  +  4Fe(OH)3  +  3O2  .  With  barium  salts  it  precipitates  a 
stable  barium  ferrate,  BaFeO4  . 

5.  Solubilities. — a. — Metal. — Iron  dissolves,  in  hydrochloric  acid  and  in  dilute 
sulphuric  acid,  to  ferrous  salts,  with  liberation  of  hydrogen  (a) ;  concentrated 
cold  H._>SO4  has  no  action,  but  if  hot,  SO2  is  evolved  and  a  ferric  salt  formed  (6); 
in  moderately  dilute  nitric  acid,  with  heat,  to  ferric  nitrate,  liberating  chiefly 
nitric  oxide  (c) ;  in  cold  dilute  nitric  acid,  forming  ferrous  nitrate  with  pro- 
duction of  ammonium  nitrate  (d),  of  nitrous  oxide  (e),  or  of  hydrogen  f) 
(Langlois,  A.  Ch.,  1856,  [3],  48,  502). 

(«)  Fe  +  H2S04  =  FeS04  +  H2 

(b)  2Fe  +  6H2S04  =  Fe2(S04)3  -f  3S02  -f  6H20 

(c)  Fe  +  4HN03  =  Fe(N03)3  +  NO  +  2H2O 

(d)  4Fe  -f  10HN03  =  4Fe(N03)2  +  NH4NO3  +  3H2O 

(e)  4Fe  +  10HN03  =  4Fe(N03)2  +  N20  +  5H2O 
(0  Fe  +  2HN03  =  Fe(N03)2  +  H2 

In  dissolving  the  iron  of  commerce  in  hydrochloric  acid,  the  carbon  which  it 
always  contains,  so  far  as  combined  in  the  carbide  of  iron,  will  pass  off  in 
gaseous  hydrocarbons  (Campbell,  Am.,  1896,  18,  836),  and  so  far  as  uncombined 
will  remain  undissolved,  as  graphitic  carbon.  The  metal  is  attacked  by  moist 
air,  forming  chiefly  2Fe2O3.3H2O  ,  iron  rust.  When  hot  iron  is  hammered,  scale 
oxide,  Fe2O3.6Fe6  ,  is  formed.  Cold  concentrated  HNO3  or  the  action 
of  the  electric  current  forms  passive  iron.  (Byers  and  Langdon,  J.  Am.  Soc., 
36,  759,  1913)  (Byers,  Ibid.,  30,  1718,  1908). 

&. — Oxides  and  hydroxides. — Ferrous  oxide  and  hydroxide  unite  with  acids 
with  rapid  increase  in  temperature,  forming  ferrous  salts,  always  mixed  with 
more  or  less  ferric  salts.  The  ferrous  salts  are  much  more  readily  prepared 
by  the  action  of  dilute  acids  upon  the  metal,  or  upon  FeC03  or  FeS  .  Fe3O4  , 
treated  with  an  insufficient  amount  of  HC1,  forms  FeCl2  and  Fe,O;!:  treated  wit  h 
HC1  sufficient  for  complete  solution,  a  mixture  of  Fed,,  and  FeCL,  is  obtained^ 
which,  when  treated  with  excess  of  ammonium  hydroxide  and  dried  at  100° 
again  exhibits  the  magnetic  properties  of  the  original.  Ferric  oxide,  Fe,O3 ,  dis- 
solves in  acids,  quite  slowly  if  the  temperature  of  preparation  of  the  oxide  has 
been  high.  Mitscherlich  (J.  pr.,  1860,  81,  110)  recommends  warm  digestion  with 
t«n  parts  of  a  mixture  of  sulphuric  acid  and  water  (8-3),  If  the  oxide  be 


156  IRON.  §126,  be. 

heated  with  alkalis  or  alkali  carbonates,  it  then  dissolves  much  more  readily  in 
acids.  Ferric  hydroxide,  Fe(OH)3 ,  is  insoluble  in  water  (for  a  soluble  colloidal 
ferric  hydroxide,  see  Sabanejeff,  C.  C.,  1891,  i,  11),  readily  soluble  in  acids  to 
ferric  salts.  Freshly  precipitated  ferric  hydroxide  readily  dissolves  in  ferric 
chloride  and  in  chromium  chloride,  not  in  aluminum  chloride.  A  solution  of 
ferric  hydroxide  in  ferric  chloride  is  soluble  in  water  after  evaporation  to  dry- 
ness  if  not  more  than  ten  parts  of  Fe,03  are  present  to  one  of  the  FeCl3  (Be- 
champ,  A.  Ch.,  1859,  (3),  56,  306) 

c. — Salts. — Ferrous  salts,  in  crystals  and  in  solution,  have  a  light  green 
color.  Solutions  of  the  salts  have  a  slight  acid  reaction  toward  litmus. 
The  sulphate  FeS04.7H20 ,  is  efflorescent;  the  chloride,  bromide,  iodide, 
and  citrate  are  deliquescent.  Solutions  of  all  ferrous  salts  are  unstable, 
gradually  changing  to  basic  ferric  salts,  more  or  less  insoluble  in  water. 
The  carbonate,  hydroxide,  phosphate,  borate,  oxalate,  cyanide,  ferro- 
cyanide,  ferricyanide,  tartrate,  and  tannate  are  insoluble  in  water. 
Ferric  salts  in  solution  have  a  brownish-yellow  color,  redden  litmus  and 
color  the  skin  yellow.  The  chloride,  bromide,  nitrate,  and  sulphate  are 
deliquescent.  The  ferrocyanide,  tannate,  borate,  phosphate,  basic  acetate, 
and  sulphite  are  insoluble  in  water;  the  sulphate  is  soluble  in  alcohol 
(separation  from  ferrous  sulphate).  Ferric  chloride  is  soluble  in  ether 
saturated  with  hydrochloric  acid,  separation  from  aluminum  (Gooch  and 
Havens,  Am.  8.,  1896,  152,  416).  Solutions  of  ferric  salts,  when  boiled, 
frequently  precipitate  a  large  portion  of  the  iron  as  basic  salt,  especially 
if  other  soluble  salts  are  present  (Fritsche,  Z.  angeir.,  1888,  227;  Pickering, 
J.  C.,  1880,  37,  807)  (§70,  5d  footnote). 

6.  Reactions,  a. — The  alkali  hydroxides  precipitate  ferrous  hydroxide. 
Fe(OH)2 ,  white  if  pure,  but  seldom  obtained  sufficiently  free  from  ferric 
hydroxide  to  be  clear  white,  and  quickly  changing,  in  the  air,  to  ferroso- 
ferric  hydroxide,  of  a  dirty-green  to  black  color,  then  to  ferric  hydroxide 
(4),  otf  a  reddish-brown  color.  The  fixed  alkalis  adhere  to  this  precipitate. 
Ammonium  chloride  or  sulphate,  sugar,  and  many  organic  acids,  to  a  slight 
extent,  dissolve  the  ferrous  hydroxide  or  prevent  its  formation  (§§116  and 
117).  The  soluble  carbonates  precipitate,  from  purely  ferrous  solutions, 
ferrous  carbonate,  FeC03 ,  white  if  pure,  but  soon  changing,  in  the  air,  to 
the  reddish-brown  ferric  hydroxide. 

Solutions  of  ferric  salts  are  precipitated  by  the  alkali  hydroxides  and 
carbonates  as  ferric  hydroxide,  Fe(OH)3 ,  variable  to  Fe203.H20 — FeO(OH)  — 
reddish-brown  insoluble  in  excess  of  the  reagents  (distinction  from  alumi- 
num and  chromium  which  are  soluble  in  excess  of  the  fixed  alkali  hy- 
droxides and  from  cobalt,  nickel  and  zinc  which  are  soluble  in  ammonium 
hydroxide).  Salts  of  the  fixed  alkalis  adhere  to. this  precipitate  with  great 
tenacity  and  the  precipitate  obtained  from  the  use  of  the  fixed  alkali 
carbonates  invariably  contains  traces  of  a  carbonate.  Freshly  precipitated 
barium  carbonate  completely  precipitates  ferric  salts  in  the  cold  as  ferric 


§126,  6*.  IRON.  15? 

hydroxide  (separation  of  ferric  iron,  with  aluminum  and  chromium,  from 
ferrous  iron,  cobalt,  nickel,  manganese,  and  zinc;  2FeCl;{  +  3BaC03  + 
3H20  =  2Fe(OH)3  +  3BaCl2  +  3C02).  The  mixture  should  be  allowed  to 
stand  several  hours  (chromium  precipitates  more  slowly  than  aluminum 
or  iron),  and,  sulphates  must  be  absent,  as  freshly  precipitated  barium 
carbonate  reacts  with  solutions  of  the  sulphates  of  the  fourth  group;  e.  g., 
NiS04  +  BaC03  =  NiC03  +  BaS04 .  The  reaction  takes  place  most  read- 
ily if  the  metals  be  present  as  chlorides.  If  the  precipitate  obtained  be 
treated  with  an  excess  of  dilute  sulphuric  acid  the  ferric  hydroxide  dis- 
solves, leaving  the  excess  of  barium  as  the  insoluble  sulphate.  Freshly 
precipitated  carbonates  of  Ca ,  Mg ,  Mn ,  Zn ,  and  Cu  react  similar  to  the 
barium  carbonate. 

I). — Oxalic  acid  and  soluble  oxalates  precipitate  from  solutions  of  ferrous 
salts,  ferrous  oxalate,  FeC2O4 ,  yellowish-white,  crystalline,  sparingly  soluble  in 
hot  water,  soluble  in  HC1  ,  HNO3  and  H,SO4;  ferric  salts  are  not  precipitated 
by  oxalates  except  as  reduction  to  ferrous  oxalate  takes  place. 

The  acetates,  as  NaC2H302 ,  form  in  solutions  of  ferric  salts  a  dull  red  * 
solution  of  ferric  aceiate,  Fe(C2H302)3 ,  which  upon  boiling  is  decomposed 
and  precipitated  as  basic  ferric  acetate  of  variable  composition  (separation 
of  iron  and  aluminum  from  phosphoric  acid  (d),  chromium,  and  the  metals 
of  the  fourth  group).  The  red  colored  ferric  acetate  solution  is  not 
decolored  by  mercuric  chloride  (distinction  from  Fe(CNS)3).  The  basic 
precipitates  are  soluble  in  HC1 ,  HN03  and  H2S04  and  are  transposed  by 
alkali  hydroxides. 

Tannic  acid  precipitates  concentrated  solutions  of  ferrous  salts:  ferric  salts 
are  precipitated  as  blue-black  ferric  tannate  (the  basis  of  common  ink),  insoluble 
in  water  or  acetic  acid,  very  soluble  in  excess  of  tannic  acid.  Ferric  salts  are 
completely  precipitated  by  ammonium  succinate  from  hot  solutions  (Young, 
J.  C.,  1880,  37,  674).  Both  ferrous  and  ferric  salts  (not  nitrates)  slightly  acid 
are  completely  precipitated  by  a  solution  of  nitroso  B.  naphthql  (separation 
from  aluminum  and  chromium)  (Knorre,  B.,  1887,  20,  283;  Menicke,  Z.  uiu.n'ir., 
1888,  5).  If  the  Fe'"  be  in  excess  of  the  PO4  the  phosphate  will  all  be  pre- 
cipitated. Hydrochloric  acid  should  be  absent,  i.  e.,  excess  of  NaC,H:!0,  should 
be  added  (Knorre,  Z.  angew.,  1893,  267). 

Potassium  cyanide  gives  with  solutions  of  ferrous  salts  a  yellowisK-red  pre- 
cipitate, which  dissolves  in  excess  of  the  reagent  to  potassium  ferrocyanide, 
K4Fe(CN)8;  with  solutions  of  ferric  salts,  ferric  hydroxide  is  precipitated  with 
evolution  of  hydrocyanic  acid  (equation  (a),  page  156). 

Potassium  ferrocyanide  precipitates  ferrous  salts  as  potassium  ferrous 
ferrocyanide  (b),  K2FeFe(CN)6 ,  (Everitt's  salt),  bluish-white,  insoluble  in 

*Meconic  acid  and  formic  acid  form  red  solutions  with  ferric-  salts  :  benzole  acid  gives  a  flesh 
colored  precipitate;  phenol,  creosote,  saligenin,  and  other  hydroxy  aromatic  derivatives  give 
a  blue  to  violet  color.  Morphine  gives  a  blue  color.  The  following  is  recommended  as  a  very 
satisfactory  test  for  a  trace  of  iron  in  copper  sulphate.  Dissolve  one  gram  of  the  CuSO4  in  five 
cc.  of  water,  add  five  cc.  of  a  ten  per  cent.  <  therial  solution  of  salicylic  acid.  Jf  the  layer  of 
eontact  assumes  a  violet  color  iron  is  present  (Grigge,  Z.,  1895,  34,  450), 


158  IRON.  §126,  6*. 

acids,  transposed  by  alkalis  (c).  This  is  converted  into  Prussian  blue 
(see  below),  gradually  by  exposure  to  the  air,  immediately  by  oxidizing 
agents  (d).  With  ferric  salts,  ferric  ferrocyanide  (e),  Fe4(Fe(CN)6)3 ,  Prus- 
sian blue,  is  formed,  insoluble  in  acids,  decomposed  by  alkalis  (/).  If  the 
reagent  be  added  in  strong  excess  the  precipitate  is  partially  dissolved  to 
a  blue  liquid.  Strong  acids  should  not  be  present  as  they  color  the  re- 
agent blue.  In  neutral  solutions  diluted  to  one  in  500,000  the  iron  may  be 
detected  (Wagner,  Z.,  1881,  20,  350).  The  ferrocyanides  are  transposed 
by  KOH  and  decomposed  by  fusion  with  NaNO;s  and  Na,C03 ,  the  iron  being 
obtained  as  Fe203  (Koningh,  Z.  angew.,  1898,  463).  Potassium  ferri- 
cyanide  precipitates  from  dilute  solutions  of  ferrous  salts  ferrous  ferri- 
cyanide  (g),  Fe3(Fe(CN)6)2  (Turnbull's  blue),  dark  blue,  insoluble  in  acids, 
transposed  by  alkali  hydroxides  (/?);  with  ferric  salts  no  precipitate  is 
obtained,  but  the  solution  is  colored  brown  or  green  (i).  This  is  a  very 
important  reagent  for  the  detection  of  the  presence  of  even  traces  of 
ferrous  salts  in  the  presence  of  ferric  salts.  As  iron  is  so  readily  oxidized 
or  reduced  by  various  reagents  the  original  solution  should  always  Be 
tested.  The  solutions  should  also  be  sufficiently  diluted  to  allow  the 
detection  of  the  precipitate  of  the  ferrous  ferricyanide  in  the  presence  of 
the  dark  colored  liquid  due  to  the  presence  of  ferric  salts.  If  no  precipi- 
tate be  obtained  (indicating  absence  of  ferrous  iron)  a  drop  of  stannous 
chloride  or  some  other  strong  reducing  agent  constitutes  a  delicate  test 
for  ferric  salts  and  reconfirms  the  previous  absence  of  ferrous  salts. 
Potassium  thiocyanate  gives  no  reaction  with  ferrous  salts;  with  ferric 
salts  the  Uood  red  ferric  thiocyanate,  Fe(CNS)3  (solution),*  is  formed  (/). 
This  constitutes  an  exceedingly  delicate  test  for  iron  in  the  ferric  condi- 
tion (the  original  solution  should  always  be  tested).  According  to  Wagner 
(Z.,  1881,  20,  350)  one  part  of  iron"  as  ferric  salt,  may  be  detected  in 
1,600,000  parts  of  water.  The  red  salt  of  ferric  thiocyanate  is  freely 
soluble  in  water,  alcohol,  and  ether;  it  is  extracted  by  ether  from  aqueous 
solutions  and  thus  concentrated,  increasing  the  delicacy  of  the  test  (Natan- 
son,  A.,  1864,  130,  246).  The  red  color  of  the  liquid  is  destroyed  by 
mercuric  chloride  (fc),  also  by  phosphates,  borates,  acetates,  oxalates,  tar- 
trates,  racemates,  malates,  citrates,  succinates,  and  the  acids  of  these  salts. 
Nitric  and  chloric  acids  give  red  color  with  potassium  thiocyanate,  re- 
moved by  heat. 

*  The  quantity  of  non-dissociated  Fe(CMS)3 ,  to  which  the  color  is  due,  is  increased  by  an  ex- 
cess of  either  of  the  products  of  the  dissociation.  The  test  for  iron  is  therefore  more  delicate 
if  considerable  KCNS  is  added.  The  decoloration  by  HgCl2  is  due  to  the  breaking-  up  of  the 
Fe(CNS)3  to  form  Hg(CNS)a  which  is  even  less  dissociated  in  water  solution  than  HgCl? , 


§126,  M.  tltott.  159 

(a)  FeCl3  +  3KCN  +  3H20  =  Fe(OH)3  +  3KC1  +  3HCBT 
(&)    PeS04  +  I^FefCN),.  =  K,FeFe(CN)0  +  K2S04 
(c)    K,FeFe(CN)0  +  2KOH  =  Fe(OH),  +  K4Fe(CN)6 

^K.FeFeCCN),  +  0,  +  ^HCl  =  Fe4(Fe(CN)0)3  +  K4Fe(CN)6  +  4KC1  +  2H20 
4FeCl3  +  3K,Fe(CN)0  — -  Fe4(Fe(CN)u)3  +  12KC1 
Fe4(Fe(CN)0)3  +  12KOH  =  4Fe(OH)3  +  3K4Fe(CN)a 
3FeS04  4-  2K3Fe(CN)(j  =Fes(Fe(CN)0),  +  3K2S04 
Fe3(Fe(CN)6)2  +  GKOH  =  3Fe(OH)2  +'2K3Fe(CN)6 
FeCl3  +  K,Fe(CN)G  =  FeFe(CN)e  +  3KC1 
FeCl3  +  3KCNS  —  Fe(CNS).,  +  3KC1 
(fc)  2Fe(CNS)3  +  3HgCl,  =  3Hg(CNS)2  +  2FeCls 

c. — Nitric  acid  readily  oxidizes  all  ferrous  salts  to  ferric  salts,  the  reac- 
tion being  hastened  by  the  aid  of  heat.  As  the  iron  is  reduced  to  the 
ferrous  condition  in  the  precipitation  of  the  metals  of  the  second  group 
with  hydrosulphuric  acid,  the  oxidation  with  nitric  acid  is  necessary  to 
insure  the  precipitation  of  all  the  iron  as  hydroxide  in  the  third  group 
(<oa  and  §117). 

d. — Hypo-phosphorous  acid  reduces  ferric  salts  to  ferrous  salts.  From 
solutions  of  ferrous  salts,  alkali  phosphates,  as  Na2HP04 ,  precipitate 
secondary  ferrous  phosphate.  FeHP04 ,  mixed  with  the  tertiary  salt, 
Fe3(P04)2 ,  white  to  bluish  white,  soluble  in  mineral  acids.  By  the  addi- 
tion of  an  alkali  acetate,  the  precipitate  consists  of  the  tertiary  phosphate 
alone:  3FeS04  +  2Na2HP04  +  2NaC2IL02  ==  Fe3(P04)2  -f  "3Na2S04  + 
2HC2H,}02 .  Ferric  salts  are  precipitated  as  ferric  phosphate,  FeP04 , 
scarcely  at  all  soluble  in  acetic  acid,  but  readily  soluble  in  hydrochloric, 
nitric  and  sulphuric  acids.*  Hence  ferric  iron  in  mineral  acid  solution 
is  precipitated  by  phosphoric  acid  with  co-operation  of  alkali  acetates : 
FeCl3  +  H3P04  +  3NaC2H302  =  =  FeP04  +  3NaCl  +  3HC2H302 .  If 
metals  of  the  fourth  group  and  the  alkaline  earths  be  present 
with  phosphoric  acid  they  are  precipitated  with  the  third  group 
metals  by  ammonium  hydroxide  in  the  usual  course  of  analysis 
(§146  and  //.)  ;  phosphates  of  Co,  Ni,  and  Zn  being  redissolved 
by  the  excess  of  ammonium  hydroxide,  To  prevent  this  general  pre- 
cipitation with  the  metals  of  the  third  group,  when  phosphates  are 
present,  the  acid  solution  (after  removal  of  the  second  group  by  hydro- 

*  Equilibrium  requires  that  a  weak  acid,  as  phosphoric,  be  present  for  the  most  part  as  the 
non-dissociated  molecule.  But  FePO4 ,  as  any  neutral  salt,  is  dissociated,  so  far  as  it  dissolve  s 
in  water,  into  its  ions,  as  is  also  the  strong  hydrochloric  acid.  Ih-ing-iu^  these  tog-ether  will  re- 
sult in  the  union  of  the  H  ion  of  the  acid  and  the  PO4  ion  to  non-dissociated  H3PO4 ,  thus 
maintaining1  the  equilibrium  for  H3PO4 ,  but  disturbing-  that  between  solid  and  dissolved 
PePO4 ,  which  requires  a  certain  concentration  of  PO4  ions.  To  restore  the  latter  more  PePO4 
dissolves,  only  to  react  with  the  H  ions  as  before,  and  this  process  continues  until  the  H  ions 
of  the  hydrochloric  acid  are  reduced  to  such  small  quantity  as  to  be  in  equilibrium  with  the 
PO4  ions  or,  if  the  HOI  is  in  excess,  until  the  FePO4  is  entirely  dissolved.  This  process  takes 
place  whenever  a  strong-  acid  dissolves  the  raH  of  a  weak  one.  It  is  analogous  to  the  solution 
of  a  base  in  an  acid,  forming  non-dissociated  water. 


160  IRON.  §126,  6*. 

gen  sulphide  and  the  expulsion  of  the  gas  by  boiling)  is  treated  with  an 
excess  of  sodium  acetate.  Acetic  acid  is  liberated  by  the  mineral  acid 
and  phosphates  of  iron,  aluminum  and  chromium  are  precipitated.  If  no 
more  phosphoric  acid  remains  in  the  solution,  the  remainder  of  the  iron, 
aluminum  and  chromium  may  be  precipitated  by  means  of  ammonia.  If 
on  the  other  hand  the  solution  contains  phosphoric  acid  it  must  be  re- 
moved. For  this  purpose  ferric  chloride  is  added  drop  by  drop,  until  a  red 
color  indicates  complete  precipitation  of  the  phosphate  and  forma- 
tion of  ferric  acetate.  The  mixture  is  then  boiled  and  filtered  hot.  All 
of  the  phosphoric  acid  present  is  thus  precipitated  and  separated  from 
the  metals  of  the  remaining  groups.  Care  should  be  taken  to  avoid  an 
excess  of  the  ferric  chloride  as  the  ferric  phosphate  is  somewhat  soluble 
in  ferric  acetate  solution.  The  alkali  hydroxides  transpose  ferric  phos- 
phate (freshly  precipitated),  forming  ferric  hydroxide  and  alkali  ph  sphate, 
The  transposition  is  not  complete  in  the  cold.  With  fixed  alkali  hydroxide 
aluminum  phosphate  is  dissolved,  thus  effecting  a  separation  from  chrom- 
ium and  iron.  Ferric  phosphate  warmed  with  ammonium  sulphide  forms 
ferrous  sulphide,  ammonium  phosphate  and  sulphur:  4FeP04  -f-  6(NH4)2S 
=  4FeS  +  4(NH4)3P04  +  S2 . 

e. — Hydrosulphuric  acid  is  without  action  upon  ferrous  salts  in  acid  or 
neutral  solutions,  except  a  slight  precipitate  is  formed  with  neutral  fer- 
rous acetate.  Alkali  sulphides  and  H2S  in  alkaline  mixture,  form  ferrous 
sulphide,  FeS ,  black,  insoluble  in  excess  of  the  reagent,  readily  soluble  in 
dilute  acids  with  evolution  of  hydrogen  sulphide.  The  moist  precipitate 
is  slowly  converted,  in  the  air,  to  ferrous  sulphate  and  finally  to  basic 
ferric  sulphate,  Fe20(S04)2 .  Ferric  salts  are  reduced  to  ferrous  salts  with 
liberation  of  sulphur  by  H2S  (1\  or  soluble  sulphides,  the  latter  at  once 
reacting  to  precipitate  ferrous  sulphide  (2) : 

(1)  4FeCl3  +  2H,S  =  4Fe€l2  +  4HC1  +  S2 

(2)  4FeCl3  +  eCNHO.JS  =  4FeS  +  12NH4C1  +  S2 

After  the  removal  of  the  metals  of  the  second  group  by  H2S ,  the  iron 
present  will  always  be  in  the  ferrous  condition  (it  will  therefore  be  neces- 
sary to  test  the  original  solution  to  find  the  condition  of  the  iron  at  the 
beginning  of  the  analysis).  The  excess  of  H2S  should  be  removed  by 
boiling  and  the  iron  oxidized  by  carefully  adding  nitric  acid  drop  by  drop 
and  boiling  until  the  solution  assumes  a  pale  straw  color  (6&).  If  this  be 
done  the  iron  will  be  completely  precipitated  in  the  third  group  by  the 
ammonium  hydroxide  (6a). 

Ferrous  sulphite  is  but  little  soluble  in  pure  water,  easily  soluble  in  excess  of 
sulphurous  acid,  to  a  colorless  solution.  The  moist  salt  oxidizes  rapidly  on 
exposure  to  the  air  (Fordos  and  Gelis,  J.  Pharm.,  1843,  (3),  4,  333).  Feme 
sulphite  is  only  known  as  a  red  solution  formed  by  the  action  of  SO,,  upon 
freshly  precipitated  Fe(OH)3,  rapidly  reduced  to  the  ferrous  condition  accord- 


§126,  8.  JRON.  161 

ing  to  the  following  equation:  Fe2(SO3)3  =  FeSO3  +  FeS2O6  (Gelis,  C.  <?.,  1862, 
896).  Ferrous  thio&ulphate,  FeS,O:!,  is  formed,  tog-ether  with  some  FeS  and  FeSO3, 
by  the  action  of  SO,  upon  Fe°  (For doe  and  Gelis,  I.  c.).  Ferric  salts  arc  reduced 
by  sodium  thiosulphate  to  ferrous  salts  in  neutral  solutions  with  formation  of 
sodium  tetrath  innate:  2FeCl,  +  2Na,S,Os  =  2FeCL  +  2NaCl  -f  Na,S4O0  (Fordos 
and  Gelis,  C.  r.,  1842,  15,  920);  in  acid  solutions  sulphuric  acid  and  sulphur  are 
formed:  4FeCl3  +  2Na,,S2O3  +  2H2O  =  4FeCL  +  4NaCl  +  zHSO4  +  S2  (Men- 
schutldn,  78).  Ferric  iron  is  precipitated  as  basic  nitrate  by  the  addition  of  a 
solution  of  ammonium  sulphate  to  a  solution  of  iron  in  HNO,,  evaporated  to 
dryness  and  taken  up  with  water  (separation  from  aluminum)  (Beilstein  and 
Luther,  C.  C.,  1891,  i,  809). 

/. — Chlorides  and  bromides  of  both  ferrous  and  ferric  iron  are  formed 
but  only  ferrous  iodide  exists.  Ferric  salts  are  reduced  to  ferrous  salts 
by  hydriodic  acid  with  liberation  of  iodine. 

g. — Soluble  arsenites  and  arsenates  precipitate  solutions  of  ferrous  and  ferric 
salts,  forming'  the  corresponding-  arsenites  and  arsenates.  Basic  ferric  arsenite, 
4Fe.O3.As.03  +  5H20  ,  is  formed  when  an  excess  of  ferric  hydroxide  is  added 
to  arsenous  acid.  It  is  insoluble  in  acetic  acid.  It  is  formed  when  moist 
ferric  hydroxide  is  given  as  an  antidote  in  case  of  arsenic  poisoning-  (§69,  GZ 
and  G'e;  D.,  3,  352). 

li.  Ferrous  salts  are  rapidly  oxidized  to  ferric  salts  by  solutions  of  chro- 
mates,  the  chromium  being  reduced  to  the  triad  condition  (9  and  10). 
With  ferric  salts  potassium  chromate  forms  a  reddish-brown  precipitate. 

i. — Zinc  oxide  precipitates  solutions  of  Fe'"  ,  Al ,  Cr"'  and  Cu  completely  and 
Pb  partially,  effecting-  a  separation  of  these  metals  from  Mn  ,  Co  and  Ni 
(Meineke,  Z.  angew.,  1888,  258). 

7.  Ignition. — The  larg-er  number  of  iron  salts  are  decomposed,  as  solids,  by 
heat;  FeCl3  vaporizes  partly  decomposed,  at  a  very  little  above  100°.  Igni- 
tion in  the  air  changes  ferrous  compounds,  and  ignition  on  charcoal  or  by 
reducing  flame  changes  ferric  compounds  to  the  magnetic  oxide,  which  is 
attracted  to  the  magnet.  Ferrous  oxalate  ignited  in  absence  of  air  gives  FeO  . 
Ferric  oxide  ignited  in  a  current  of  hydrogen  gives  Fe304  from  330°  to  440°,  FeO 
from  500°  to  600°,  and  Fe°  above  600°  (Moissan,  A.  Ch.,  1880,  (5),  21,  199). 

In  the  outer  flame,  the  borax  bead,  when  moderately  saturated  with  any 
compound  of  iron,,  acquires  a  reddish  color  while  hot,  fading  and  becoming 
light  yell  wo  when  cold,  or  colorless,  if  feebly  saturated.  The  same  bead,  held 
persistently  in  the  reducing  flame,  becomes  colorless  unless  strongly  saturated, 
when  it  shows  the  pale  green  color  of  ferrous  compounds.  The  reactions  with 
microcosmic  salt  are  less  distinct,  but  similar.  Cobalt,  nickel,  chromium  and 
copper  conceal  the  reaction  of  iron  in  the  bead. 

Ferric  compounds,  heated  briefly  in  a  blue  borax  bead  holding  a  very  little 
cupric  oxide,  leave  Jthe  bead  blue;  ferrous  compounds  so  treated  change  the 
blue  bead  to  red — the  color  of  cuprous  oxide. 

8.  Detection. — After  removal  of  the  first  two  groups  the  iron  (now  in 
the  ferrous  condition)  is  oxidized  by  HN03  and  then  precipitated  in  pres- 
ence of  NH4C1  with  Al  and  Cr'"  by  an  excess  of  NH4OH  .  The  Al  is  re- 
moved by  boiling  with  excess  of  KOH  .  If  more  than  traces  of  Fe  be 
present  it  is  detected  in  presence  of  the  Cr(OH)3 ,  by  dissolving  in  HC1 
and  obtaining  the  blood-red  solution  with  KCNS  .  In  case  Cr  be  present 
in  great  excess  the  Cr(OH)3  and  Fe(OH)3  are  fused  on  a  platinum  foil  with 


162  IRON.  §126,  9,  (1). 

Na2C03  and  KNOn ,  oxidizing  the  Cr  to  a  chromate  soluble  in  water.  After 
filtering,  the  precipitate  of  Fe203  is  dissolved  in  HC1  and  tested  with  KCNS. 
The  original  solution  must  be  tested  to  determine  whether  the  iron  was 
present  in  the  ferrous  or  ferric  condition.  A  portion  of  the  original 
solution  acidified  with  HC1  gives  blood  red  color  with  KCNS  if  Fe'"  is 
present,  no  color  for  the  Fe".  Another  portion  gives  a  blue  precipitate  with 
K3Fe(CN)G  if  Fe"  is  present,  only  a  brown  or  green  color  for  the  Fe'"(C#). 
Ferrous  iron,  in  the  presence  of  ferric,  may  be  detected  as  follows :  To 
a  small  portion  of  the  solution  to  be  tested  an  equal  volume  of  concentrated 
sulphuric  acid  is  added  without  mixing.  After  cooling  the  solution  a 
a  crystal  of  potassium  nitrate  is  added.  On  shaking  the  test  tube  gently  a 
red  to  brown  color  develops  if  ferrous  iron  is  present.  Blum,  Z.,  44,  10. 

9.  Estimation. — (1)  After  oxidation  to  Fe'"  ,  if  necessary,  it  is  precipitated 
with  N~H,OH  ,  dried,  ignited  to  a  dull-red  heat  and  weighed  as  Fe2O3  .  (2)  By 
precipitation  with  nitroso-/3-naphthol  in  slightly  acid  solution  (Knorre,  B.,  L887, 
20,  283).  Volumetrically:  (3)  As  ferrous  iron,  by  titration  with  a  standard 
solution  of  KMn04:  10FeS04  +  2KMn04  +  8H2SO4  =  f>Fe,(S04)3  +  K,SO4  + 
2MnSO4  +  8H_,0  .  (//)  By  titration  with  a  standard  solution  of  K,Cr,O7  ,  using 
a  solution  of  K3Fe(CN)G  as  an  external  indicator:  6FeSO4  +  K.CrA  +  7H,SO4  = 
3Fe2(S04)3  +  K,S04  +  Cr2(SO4)3  +  7H2O  .  (o)  As  ferric  iron,  by  titration  with 
a  standard  solution  of  Na,S,03 ,  using  KCNS  as  an  indicator:  2FeCl3  +  2Na,S,03 
=  2FeCl>  +  NaJS^Oe  +  2NaCl .  A  few  drops  of  a  solution  of  CuSO4  are  added, 
which  seems  to  hasten  the  reaction  and  gives  more  accurate  results;  or  use 
excess  of  the  N"a2SoO3  and  titrate  back  with  standard  iodine  (Crafts,  J.  €.,  1873, 
26,  1162).  (6)  The  iron  as  ferric  salt  is  treated  with  an  excess  of  a  standard 
SnCl,  solution,  the  excess  of  the  SnCl,  being  determined  by  a  standard  solution 
of  iodine  in  potassium  iodide:  2FeCl3  +  SnCL  =  2FeCl2  +  SnCl4 .  (7)  Potas- 
sium iodide  is  added  to  the  nearly  neutral  ferric  chloride;  the  flask  is  stoppered 
and  warmed  to  40°.  The  iodine  set  free  is  titrated  by  standard  Na-,S203 
(very  accurate  for  small  amounts  of  iron).  (8)  When  present  in  traces  it  is 
determined  colorimetrically  as  Fe(CNS)3  in  etherial  solution  (Lunge,  Z.  angeiv., 
1894,  669). 

10.  Oxidation. — Metallic  iron  precipitates  the  free  metals  from  solu- 
tions of  An  ,  Pt ,  Ag  ,  Hg ,  Bi ,  and  Cu  (separation  from  Cd). 

Solutions  of  Fe"  are  changed  to  Fe"'  solutions  by  treating  with  solutions 
of  Au ,  Ag ,  CrVI,  Mnvn,  Mnvl,  and  H202 .  In  presence  of  some  dilute 
acid,  such  as  H2S04  or  H3P04  by  Pb02 ,  Pb304 ,  Mn304 ,  Mn02 ,  Mn203 , 
Co203 ,  Ni203 .  The  following  acids  also  oxidize  Fe"  to  Fe'",  HN02 ,  HN03 , 
HC10 ,  HC102 ,  HC10,  ,  ILS04  (if  concentrated  and  hot),  HBrO ,  HBr03 
HIO., ,  also  Br ,  Cl .  Br  and  Cl  in  presence  of  KOH  changes  Fe"  and  Fe"' 
to  K2Fe04 .  Barium  ferrate  is  the  most  stable  of  the  ferrates ;  they  are 
strong  oxidizers,  acting  upon  nitrites,  tartrates,  glycerol,  alcohol,  ether* 
ammonia,  etc.  (Resell,  J.  Am.  Soc.,  1895,  17,  760). 

Fe'"  is  reduced  to  Fe"  by  solutions  of  Sn",  Cu',  H3P02 ,  H3PO, ,  H2S , 
H2S03 ,  Na2S203 ,  and  HI .  Also  by  nascent  hydrogen,  or  by  any  of  the 
metals  which  produce  hydrogen  when  treated  with  acids,  including  Pb , 
As  ,  Sb  ,  Sn  ,  Bi ,  Cu  .*,  Cd ,  Fe  ,  Al ,  Co  ,  Ni ,  Zn  ,  and  Mg  f. 

*  Carnegie,  J.  C.,  1888,  53,  468.  t  Warren,  C.  N.,  1889,  60, 187, 


§127. 


ANALYSIS    OF   THE  IRON  GROUP. 


163 


§127.  .TABLE  FOR  ANALYSIS  or  THE  IRON  OR  THIRD  GROUP  (Phosphates 
and  Oxalates  being  absent).     See  §312. 

To  the  clear  filtrate  from  the  Second  Group,  in  which  H2S  will  cause  no  pre- 
cipitate (§80),  and  freed  from  H2S  by  boiling,  add  a  few  drops  of  Nitric 
Acid  and  boil  an  instant  (to  oxidize  the  ferrous  iron  *).  Immediately  add 
Ammonium  Chloride  (§134,  56;  §189,  56)  and  an  excess  (§136,  6a)  of 
Ammonium  Hydroxide  (§116).  If  there  is  a  precipitate,  filter  and  wash. 

Precipitate:  A1(OH)3|  Cr(OH)3,  Fe(OH)3  [Mn(OH),] . 

Pierce  the  point  of  the  filter,  and  with  a  little  water  wash  the  precipitate 
into  a  casserole  or  evaporating  dish.  Add  5  to  10  cc.  of  NaOH  and  10 
to  15  cc.  of  H2O2  or  add  a  gram  or  two  of  sodium  peroxide.  Boil  several  minutes, 
dilute  and  filter. 


Residue:  Fe    (OH)3, 
Mn(OH)3 

Dissolve  in  dilute  H2SO4, 
adding  a  few  drops  of 

Filtrate:   Na2CrO4,  NaAlO2 

Acidify  with   acetic  acid  and   divide  into  two  parts. 

H2O2  if  necessary.  Di- 

(1)  Test    for     chromium. 

(2)   Test    for     aluminum. 

vide  in  two  parts. 

Add  a  few  drops  of  lead 

Add  (NH4)2CO3  in  ex- 

(1) Add  KCNS.  Blood- 
red  color  indicates  iron 

/s  -tno  njLN 

acetate.       The     yellow 
lead      chromate      indi- 
cates   chromium    (§  57, 

cess  and  heat.  A  pre- 
cipitate is  A1(OH)3. 

(,§  1^0,  DO). 

6/i). 

The    same    result    is    ob- 

(2) Boil  with  HNO3    and 

tained  with  nearly 

PbO2  or  PbjO4.  Red- 
dish purple  solution  in- 
dicates manganese. 

Iron  being  found,  to  de- 
termine whether  it  is 
ferric  or  ferrous,  or 
bothf,  in  the  original 
solution,  test  the  latter, 
after  acidulating  with 
hydrochloric  acid,  with 
KCNS  for  ferricum, 
and  with  K3Fe(CN)6 
forferrosum  (§  126,  66). 

If    the    original    solution 
contains  a  chromate  it 
will  be  yellow   (normal 
chromate),  or  red  (acid 
chromate),      and      will 
give    the    reactions   for 
chromates  with 
Pb(C2H;i02)2  ,       BaCl2, 
etc.  (§125,  6h).     If  the 
chromium  is  present  as 
achromic  salt,Cr2(SO4)3, 
the   solution   will   have 
a  green  or  bluish-green 
color  and  will  give  the 

equal  certainty  by  add- 
ing an  excess  of  NH4C1 
to  the  alkaline  solution 
§124,  6a;  §130). 

Lead  and  antimony  give 
similar  results  if 
(through  carelessness) 
they  have  not  been 
removed  (§  131,  6). 

general  reactions  as  de- 

scribed at  §  125,  6. 

Chromates  should  be  re- 

duced  by   boiling  with 
HC1   and   C2H5OH    be- 

fore    proceeding     with 

the    regular    course    of 

analysis    (§  125,  6/). 

Study  §  136,  §  128,  §  129, 
§  130,  and  §131. 

Study   §136,    §128,    §129, 
§130,  §131. 

Study  §  136,  §  128,  §129. 
§  131,  6  and  §  124,  6. 

*  In  the  filtrate  from  the  Second  Group  iron  is  necessarily  In  the  ferrous  condition  (126,  6e). 
t  Ferrous  salts,  which  have  been  kept  in  the  air,  are  never  wholly  free  from  ferric  compounds. 


164  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES.  §128. 

DIRECTIONS  FOR  THE  ANALYSIS  OF  THE  METALS  OF  THE  THIRD  GROUP. 

§128.  Manipulation. — Boil  the  filtrate  from  the  second  group  (§80)  to 
expel  the  H2S  and  then  oxidize  any  ferrous  iron  that  may  be  present  by 
the  addition  of  a  few  drops  of  HN03 ,  continuing  the  boiling  to  a  clear 
straw-colored  solution  (§126,  6c): 

3FeS04  +  4HN03  =  Fe2(SO4)3  +  Fe(N03)3  +  NO  +  2H20 
Add  to  the  solution  about  one-half  its  volume  of  NH4C1  (56,  §§134  and 
189)  and  warm  and  then  add  NH4OH  in  a  decided  excess  (§135,  6a): 
MgCl2  +  NH4C1  +  NH4OH  =  NH4MgCl3  +  NH4OH 
Fe2(S04)3  +  6NH4OH  =  2Fe(OH),  +  3(NH4)2SO4 
ZnS04  +  4NH4OH  =  (NH4)2Zn02  +  (NH4)2S04  +  2H2O 

Heat  nearly  to  boiling  for  a  moment,  filter,  and  wash  with  hot  water. 
Notice  that  the  filtrate  has  a  strong  odor  of  ammonium  hydroxide  and 
set  aside  to  be  tested  for  the  metals  of  the  succeeding  groups  (§138). 

§129.  Notes.— (1)  If  the  H2S  is  not  all  expelled,  it  becomes  oxidized  by  the 
HN03  with  deposition  of  a  milky  precipitate  of  sulphur  (§257,  6#),  which 
tends  to  obscure  the  reactions  following:  6H2S  -4-  4HNO3  =  3S2  -4-  4NO  +  8H2O. 
Also  any  H2S  not  decomposed  by  the  HNO3  would  cause  a  precipitate  of  the 
sulphides  of  the  fourth  group  upon  the  addition  of  the  NH4OH:  H2S  -4-  NiCl2  + 
2NH4OH  =  NiS  +  2NH4C1  +  2H20  . 

(2)  Any  iron  that  may  have  been  present  in  the  original   solution   in   the 
ferric   condition  is  reduced  to   the  ferrous  condition   by   the   H,S    (§126,   6e) : 
4FeCl3  +  2H2S  =  4FeCl2  +  S2  +  4HC1 .     The  ferrous  hydroxide  is  not  com- 
pletely insoluble  in  the  ammonium  salts  present  (§117),  and  hence  unless  the 
oxidation  with  the  HNO,  be  complete,  some  of  the  iron  will  be  found  in  the 
next  group. 

(3)  If  considerable  iron  be  present  the  solution  becomes  nearly  black  upon 
addition  of  nitric  acid,  due  to  the  combination  of  the  nitric  oxide  with  the 
ferrous  iron  (§241,  8a).     Therefore  the  boiling,  and  addition  of  HNO3  ,  a  drop 
or  two  at  a  time,  must  be  continued  until  the  solution  assumes  a  bright  straw 
color. 

(4)  If  nitric  acid  be  added  in  excess  there  is  danger  that  Mn  will  be  oxid- 
ized to  the  triad  or  tetrad  condition  then  it  is  precipitated  with  iron  in  the 
third  group  (§134,  6a).     The  careful  addition  of  the  nitric  acid    (avoiding  an 
excess)  prevents  this  oxidation  of  the  manganese. 

(J)  Ammonium  hydroxide  precipitates  a  portion  of  Mn  (§134,  6«)  and  Mg 
(§189,  60),  but  these  hydroxides  are  soluble  in  NH4C1  (5c,  §§134  and  189): 
hence  if  that  reagent  be  added  in  excess  the  Mg  k  not  at  all  precipitated  by  the 
NH4OH,  but  the  manganese  is  oxidized  to  Mn(OH)a  and  precipitated  more  or 
less  completely  (§  134,  6a): 

2MnCl2  +  2NH4OH  =  Mn(OH)2  +  (NH4)2MnCl4 
Mn(OH)2  +  4NH4C1  =  (NH4)2MnCl4  +  2NH4OH 
2(NH4)2MnCl4  +  4NH4OH  +  O  +  H2O  =  2Mn(OH)3  +  8NH4C1 
2MgCl2  +  2NH4OH  =  Mg(OH)2  +  NH4MgCl3  +  HN4C1 
Mg(OH)2  +  3NH4C1  =  NH4MgCl3  +  2NH4OH 

Barium,  calcium,  and  strontium  are  also  precipitated  as  carbonate  because  the 
alkaline  solution  absorbs  carbon  dioxide  from  the  air.  As  much  as  15  mg.  of 
barium  may  be  present  in  this  precipitate.  (Curtman  and  Frankel,  J.  Am.  Soc., 
33,  724,  1911.)  For  detection  of  the  barium  see  §  186,  8. 

(6)  Ammonium  chloride  lessens  the  solubility  of  A1(OH)3  in  the  NH4OH 
solution  .and  effects  &n  almost  quantitative  precipitation  of  that  metal  (§117), 


§130.  DIRECTIONS  FOR   ANALYSIS  WITH  NOTES.  165 

(7)  NH4OH  precipitates  solutions  of  Co  ,    Ni  and  Zn  ,  but  these  precipitates 
are  readily  soluble  in  an  excess  of  the  NH4OH  (§116).      To  insure  the  presence 
of  an  excess  of  NH4OH  the  odor  should  be  noted  after  shaking  the  test  tube   and 
after  the  solution  has  been  heated. 

(8)  The  precipitates  of  the  hydroxides  of  Al  ,  Cr  and  Fe'"  filter  much  more 
rapidly  if  the  precipitation  takes  place  from  a  hot  solution  (§124,  4  and  6a). 

(9)  In  the  presence  of  chromium  the  filtrate  from  the  third  group  is  usually 
of  a  slight  violet  color,  due  to  the  solution  of  a  trace  of  chromium  hydroxide 
in  the  NH4OH  (§125,  6a).     Boiling-  the  solution  to  remove  excess  of  ammonia 
prevents  this. 

(10)  A  small  portion  of  the  filtrate  of  the  second  group  after  the  removal  of 
the  H2S  by  boiling  should  be  tested  for  the  presence  of  phosphates  by   am- 
monium  molybdate    (§75,   6d).     If    phosphates   are   found    to   be    present,    the 
method  of  analysis  of  the  succeeding  groups  must  be  considerably  modified. 
These  modifications  are  fully  discussed  under  §145  to  §153. 

§130.  Manipulation. — The  well  washed  precipitates  of  Al,  Cr ,  Fe'" 
and  Mn'"  hydroxides  are  transferred  to  a  small  casserole  or  evaporating 
dish  by  piercing  the  point  of  the  filter  and  washing  the  precipitate  from 
the  filter  with  as  small  a,n  amount  of  water  as  possible. 

Add  NaOH  and  H202  and  boil  for  a  minute  or  two  or  add  Na202  in 
small  portions  and  boil  until  the  chromium  is  completely  oxidized  as 
indicated  by  the  yellow  color. 

2Cr(OH)3  +  3Na2O2  =  2Na>Cr(>4  +  2H2O  +  2NaOH 
A1(OH)3  +  NaOH  =  NaAlO2  +  2H,O 

The  alkaline  solution  is  filtered  (the  filtrate  is  reserved  to  be  tested  for 
aluminum)  and  the  remaining  precipitate  dissolved  in  dilute  HC1  or 
H2S04  with  the  addition  of  a  few  drops  of  H202  if  necessary  to  dissolve 
the  manganese  hydroxide. 

Fe2O3  +  3H2  O4  =  Fe2(SO4)8  +  3H2O 

2Mn(OH)3  +  2H2SO4  +  H2O2  =  2MnSO4  +  6H2O  +  O2 

The  iron  is  tested  for  by  adding  a  few  drops  of  ammonium  or  potassium 
sulphocyanate : 

Fe2(SO4)3  +  6KCNS  =  2Fe(CNS)s  +  3K2SO4 

If  iron  has  been  found  to  be  present,  the  original  solution  acidulated 
with  HC1  (or  a  few  drops  of  the  filtrate  from  the  first  group)  should  be 
tested  with  KCNS  for  the  presence  of  ferric  iron  (§126,  66)  and  with 
K3Fe(CN)6  for  the  dark  blue  precipitate  of  Fe3[Fe(CN)6]2  indicating 
the  presence  of  ferrous  iron  (§126,  66)  : 

3FeSO4  +  2K3Fe(CN)6  -  Fe3[Fe(CN)6]2  4-  3K2SO4 
Manganese  is  tested  for  by  boiling  a  portion  of  the  solution  with  nitric 


166  DIRECTIONS  FOR   ANALYSIS  WITH  NOTES.  §131,  L 

acid  and  Pb02  or  Pb304  .*  If  manganese  is  present  a  solution  of  the 
reddish  purple  permanganic  acid  is  obtained. 

2MnSO4  +  5PbO2  +  6HNOS  =  2HMnO  4  +  2PbSO4  +  3Pb(NO3)2  +  2H20 
The  alkaline  filtrate  obtained  after  boiling  the  precipitated  hydroxides 
with  NaOH  and  H202  is  acidified  with  acetic  acid  and  one  portion  tested 
for  chromium  with  Pb(C2H302)2  a  yellow  precipitate  indicating  chromium. 

Na2CrO4  +  Pb(C2K3O2)2  =  PbCrO4  +  2NaC2H3O2  (§57,  6h) 

Another  portion  is  tested  for  aluminum  by  adding  (NH4)2C03  in  excess 
and  warming: 

NaA102  +  4HC2H302  =  A1(C2H3O2)3  +  NaC2H3O2 

2A1(C2H3O2)3  -f  3(NH4)2CO3  +  3H2O  =  2A1(OH)3  +  6NH4(C2H3O2)  +  3CO2 

a  white  gelatinous  precipitate  being  evidence  of  the  presence  of  aluminum. 
A  similar  precipitate  may  also  be  obtained  by  adding  excess  of  NH4C1 
to  the  alkaline  solution: 

2KA1O2  +  2NH4C1  +  H2O  =  A12O(OH)4  +  2KC1  +  2NH3  (§124,  6a) 

§  131.  Notes.  —  (1)  Chromium  hydroxide  is  oxidized  by  an  alkaline  oxidizing 
agent  to  a  chromate.  Sodium  peroxide  may  be  used  or  hydrogen  peroxide  and 
caustic  soda  which  unite  to  form  sodium  peroxide.  The  same  result  may  be 
obtained  by  fusing  chromium  hydroxide  on  a  platinum  foil  with  a  mixture  of  equal 
parts  of  KNO3  and 


2Cr(OH)3  +  3KN03  +  2Na2CO3  =  2Na2CrO4  +  3KNO2  +  2CO2  +  3H2O  . 

Aluminum  is  converted  into  sodium  aluminate 

2A1(OH)3  +  Na2CO3  =  2NaAlO2  +  3H2O  +  CO2  . 

(#)  Unless  the  precipitate  of  the  hydroxides  is  a  very  dark  green,  due  to  the 
presence  of  a  large  amount  of  chromium,  a  portion  of  the  precipitate  should 
be  dissolved  in  HC1  and  tested  with  KCNS  for  the  presence  of  iron.  The 
presence  of  a  moderate  amount  of  chromium  does  not  interfere. 

(3)  In  the  absence  of  chromium  the   presence  of  more  than  traces  of  iron 
gives  a  brown  color  to  the  ammonium  hydroxide  precipitate   (§126,   6a),  alu- 
minum hydroxide  being  a  white  gelatinous  precipitate. 

(4)  Manganese  remains  undissolved  when  the    mixed    hydroxides    are  treated 
with  sodium  peroxide.     When  they  are  fused  with  KNOs  and  Na^COs  the  mangan- 
ese is  converted  into  the  green  manganate.      (§  134,  7).     This  dark  green  color 
is  an  excellent  test  for  manganese  which  is  almost  invariably  present  with  the 
hydroxides  if  it  is  in  the  solution  of  the  unknown  (§  134,  6a).     On  treating  the 
fused  mass  with  water  a  green  solution  is  obtained  which  turns  purple  on  stand- 
ing or  cautiously  acidifying  with  acetic  acid. 

(«)      K2Mn04  +  8HC1  =  MnCl2  +  2KC1  +  2C12  +  4H2O 

(6)     2K2CrO4  +  10HC1  +  30^0  =  2CrCl3  +  4KC1  +  3C2H4O  +  8H2O 

(5}  The  presence  of  chromium  as  chromic  salts  is  usually  indicated  by  the 
green  or  bluish-green  color  of  the  original  solution.  Chromium  as  chromates 
(red  or  yellow)  should  be  reduced  to  chromic  salts  by  bailing  with  HC1  and 

*  As  PbOz  and  PbsCh  frequently  contain  traces  of  manganese,  the  samples  used  when  testing 
for  manganese  must  first  be  tested  as  follows:  Boil  the  oxide  with  dilute  nitric  acid  and  allow 
to  settle.  If  the  liquid  is  red,  manganese  is  present.  A  pure  sample  of  oxide  must  be  obtained 
or  the  manganese  removed  by  repeated  boiling  with  nitric  acid. 


§132,  4.  COBALT.  167 

C2H6O  before  proceeding  with  the  regular  group  separations  (§  126,  6e  and  O 
H2S  will  effect  this  reduction  but  gives  also  ;i  precipitate  of  sulphur  which 
should  be  avoided  when  convenient  to  do  so:  :2K  Cr  O7  -f-  lOHCl  +  r'H,S  = 
<iCrCl:s  +  -1KC1  +  3S2  +  MH,0  . 

(6)  Too  much  stress  cannot   be  laid  upon  1he   neerssily  for  removing  all  the 
metals   of   one   group   before   testing   the    filtrate    for   the   metals    of    1he    next 
succeeding-  group.     If  through  lack  of  sufficient  H,S  or  too  much  HC1  ,  lead  or 
antimony  are  not  completely  removed   in   the   second   group,  they   will   give   all 
the  reactions  for  aluminum  (§57,  0(7,  and   §70,  fir/);  hence  as  a  safeguard   it  is 
advised    to    test    the    white    precipitate,    indicating    aluminum,    with    HaS  .     A 
black  or  orange'  precipitate  is  evidence  of  unsatisfactory  work  and  the  student 
should  repeat  his  analysis. 

(7)  The  presence  of  a  trace  of  white  precipitate  in  the  final  test  for  aluminum 
may  he  due  to  the  presence  of  that  metal  in  the  fixed  alkali  (§124.  fir/,  footnote), 
or   it  may  be  caused  by  the  use  of  too  concentrated  fixed  alkali,  which   may 
dissolve   silica   from  the  glass  of  the  test  tubes  or  remove  it  from   the  filter 

a  per  (§249,5).  The  caustic  soda  solution  maybe  tested  by  acidifying  10  c  c 
with  HuJ,  adding  excess  of  (NH4)2CO3  and  warming.  Note  the  amount  of  reci  i- 
tate  formed  and  do  not  re  ort  aluminium  unless  a  heavier  \  recipitate  is  obtained  in 
t;  e  analysis. 

THE  ZINC  GROUP  (FOUKTH  GROUP). 

Cobalt,  Nickel,  Manganese,  and  Zinc. 

§132.  Cobalt.     Co  =  58.97.     Usual  valence  two  and  three. 

1.  Properties. — Specific  gravity,  powder  from  the  oxide  reduced  by  hydrogen, 
mean  of  five  samples,   8.957  (Rammelsberg,  Pogg.,   1849,   78,    93);    melting   foivt, 
1480°  (Cir.  B.  S.,  36,  1915).     Cobalt  is  similar  to  iron  in  appearance,  is  harder 
than  Fe  or  Ni  .     It  is  malleable,  very  ductile  and  most  tenacious  of  any  metal, 
the  wire  being  about  twice  as  strong  as  iron  wire  (Deville,  A.  Ch.,  1856,  (3),  46, 
202).     The  fine  powder  oxidizes  in  the  air  quite  rapidly  and  may  even  take  fire 
spontaneously;    in  a  compact  mass  it  is  but  little  tarnished  in  moist  air.     At  a 
white  heat  it  burns  rapidly  to  Co3O4.     It  is  attracted  by  the  magnet  and  can 
be  made  magnetic,  retaining  (unlike  steel)  its  magnetism  at  a  white  heat. 

2.  Occurrence — Cobalt   does   not   occur  in   a   free   state,    except  in   meteoric 
iron.     It  is  found  in  linnaeite  (Co,<S4);    skutterudite  (CoAs3);    smaltite  (CoAs2); 
cobaltite  (CoAsS);    wad,  asbolite  (impure  hydrous  oxides  of  Mn  ,  Fe  ,  Cu  and 
Co) ;   and  in  a  number  of  other  rare  minerals. 

3.  Preparation. — (1)    By  electrolysis  of   the  chloride.     (2)   By  heating-  with 
potassium   or  sodium.     (3)   By  heating-  any  of  the  oxides,   hydroxides   or  the 
chloride  in  hydrogen  gas.     (//)  By  fusion  of  the  oxalate  under  powdered  glass. 
(.•>)  Also  reduced  by  carbon  in  various  ways. 

4.  Oxides  and   Hydroxides.— Cobaltous  oxide,   CoO  ,   is  made    (1)    by  heating 
any  of  its  oxides  or  hydroxides  in  hydrogen  to  (not  above)  .350°;  (2)  by  ignition 
of  Co(OH)2  or  CoCO3  ,  air  being  excluded;  (3)  by  heating  Co3O4  to  redness  in 
a,  stream  of  CO,   (Russell,  J.  C.,  1863,  16,  51);  (//)  by  heating  any  of  the  higher 
oxides  to  a   white  heat  (Moissan,  A.  Ch.,  1880,  (5),  21,  242).     CobaltoilS  lii/<lro.ri<lc 
is   made   from  cobaltous  salts    by   precipitation  with   fixed   alkalis:   oxidizes   if 
exposed  to  the  air  (6«).     The  most  stable  oxide  is  the  coital toxo-cobfiltie  (Co3O4) 
tricofxill  Icfrn.rhic;  it  is  made  by  heating  any  of  the  oxides  or  hydroxides,  the 
carbonate,  oxalate  or  nitrate  to  a  dull-red  heat  in  the  air  or  in  oxygen  gas. 
Several  oxide-hydroxides  are  known,  e.g.,  Co3O2(OH)4,  Co3O(OH)r>,  Co3O3(OH)o. 
Cohdltic  o.rifJc,  Co203  ,  is  made  by  heating*  the  nitrate  just  hot  enough  for  de- 
composition, but  not  hot  enough  to  form  Co3O4  .     CobaltlC  lii/<]r<>.rid<\  Co(OH)3  , 
is  made  by  treating  any  cobaltous  salt  with  Cl  ,  HC1O  ,  Br  or  I  in  presence  of 
a  fixed  alkali  or  alkali  carbonate.     It  dissolves  in  HC1  with  evolution  of  chlo- 
rine, in  H  SO,  with  evolution  of  oxygen,  forming  a  cobaltous  salt.     CoO2  has 
not  yet  been  isolated,  but  McConnell  and  Hanes   (J.   C.,   1897,  71,  584)   have 
-ho\\u  that  it  exists  as  HaCoO,  and  in  certain  cobaltitee. 


168  COBALT.  §132,  5rt. 

5.  Solubilities. — a. — Metal. — Slowly   soluble    on   warming    in    dilute    HC1    or 
H2SO4  ,  more  rapidly  in  HNO3  ,  not  oxidized  on  exposure  to  the  air  or  when 
heated    in    contact    with    alkalis.     Like    iron,    it    may    exist    in    a   passive    form 
(Nickles,  J.  pr.,  1854,  61,  168;  St.  Edme,  C.  r.,  1889,  109,  304).     With  the  halogens 
it  forms  cobaltous  compounds   (Hartley,  J.   (7.,   1874,   27,   501).     6. — Oxides  and 
7«i/dro,rM7es.— Cobaltous    oxide    (gray-green)    and    hydroxide    (rose-red)    are    in- 
soluble in  water;  soluble  in  acids,  in  ammonium  hydroxide,  and  in  concentrated 
solutions  of  the  fixed  alkalis  when  heated    (Zimmerman,   A.,   1886,   232,   324); 
the  various  higher  oxides,  and  hydroxides  are  insoluble  in  ammonium  hydroxide 
or  chloride   (separation  from  nickelous  hydroxide   after  treating  with  iodine 
in  alkaline  mixture)   (Donath,  Z.,  1881,  20*  386),  and  are  decomposed  by  acids, 
evolving  oxygen  with  non-reducing  acids,  or  a  halogen  from  the  halogen  acids, 
and  forming  cobaltous  salts.   Co3O4  is  said  to  be  soluble  in  acids  with  great  diffi- 
culty (Gibbs  and  Geuth,  Am.  8.,  1857,  (2),  23,  257).     c—Snlts.— Cobalt  forms  two 
classes  of  salts:  coMltous,  derived  from  CoO  ,   and  cobaltic,  from   Co2O3  .     The 
latter  salts  are  quite  unstable,   decomposing  in  most  cases   at   ordinary  tem- 
peratures,  forming  cobaltous   salts.     The   cobaltous   salts   show   a    remarkable 
variation  of  color.     The  crystallized  salts  with  their  water  of  crystallization 
are  pink;  the  anhydrous  salts  are  lilac-blue.     In  dilute  solution  the  salts  are 
pink,  but  most  of  them  are  blue  when  concentrated  or  in  presence  of  strong 
acid.     A  dilute   solution   of   the   chloride  spreads   colorless  upon   white   paper, 
turning  blue  upon  heating  and  colorless  again  upon  cooling,  used  as  "  sympa- 
thetic ink." 

Cobaltous  nitrate  and  acetate  are  deliquescent;  chloride,  hygroscopic;  sulphate, 
efflorescent.  The  chloride  vaporizes,  undecomposed,  at  a  high  temperature. 

The  carbonate,  sulphide,  phosphate,  borate,  oxalate,  cyanide,  ferrocyanide 
and  ferricyanide  are  insoluble  in  water.  The  potassium-cobaltous  oxide  is  in- 
soluble; the  ammonio-cobaltous  oxide,  and  the  double  cyanides  of  cobalt  and  the 
alkali  metals,  soluble  in  water.  Alcohol  dissolves  the  chloride  and  nitrate; 
ether  dissolves  the  chloride,  sparingly,  more  so  if  the  ether  be  saturated  with 
HC1  gas  (separation  from  Ni)  (Pinerua,  C.  r.,  1897,  124,  862).  Most  of  the 
salts  insoluble  in  water  form  soluble  compounds  with  ammonium  hydroxide. 

6.  Reactions,     a. — The  fixed  alkali  hydroxides  precipitate,  from  solu- 
tions of  cobaltous  salts,  blue  basic  salts,  which  absorb  oxygen  from  the  air 
and  turn  olive  green,  as  cobaltoso-cobaltic  hydroxide;  or  if  boiled  before 
oxidation  in  the  air,  become  rose-red,  as  cobaltous  hydroxide,  Co(OH)2 . 
The  cobaltous  hydroxide  is  not  soluble  in  excess  of  the  reagent,  but  is 
somewhat  soluble  in  a  hot  concentrated  solution  of  KOH  or  NaOH   (dis- 
tinction from  Ni)  Reichel,  Z.,  1880,  19,  468).     The  cold  solution  is  blue, 
the  color  being  more  intense  when  hot  and  especially  if  a  little  glycerine 
is  present.     On  standing  it  turns  green,  then  red,  the  changes  being  more 
rapid  if  a  little  hydrogen  peroxide  solution  is  added.     Copper  also  gives 
a  blue  solution  but  does  not  give  the  other  color  changes.     Donath,  Z., 
40,     137.      Freshly    precipitated    Pb(OH)2?    Zn(OH)2,    and    HgO    pre- 
cipitate   Co  (OH)  2    from    solutions    of   cobaltous    salts   at    100°.      Ammo- 
nium   hydroxide    causes    the    same    precipitate    as    the    fixed    alkalis; 
ncomplete,    even  at  first,   because  of  the  ammonium  snlt  formed  in  the 
reaction,  and  soluble   in   excess  of  the  reagent  to  a  solution  which  turns 
brown  in  the  air  by  combination  with  oxygen,  and  is  not  precipitated  by 
potassium  hydroxide.     The  reaction  of  the  precipitate  with  ammonium 
salts  forms  soluble  double  salts  (as  with  magnesium) ;  the  reaction  of  the 


8132,  6J.  COBALT.  169 

precipitate  with  ammonium  hydroxide  produces,  in  different  conditions, 
different  soluble  compounds  noted  for  their  bright  colors,  as  (NH3)4CoCl2 , 
(NH3)nCoCl2 ,  (NHs)4CoCl3 ,  etc. 

Alkali  carbonates  precipitate  cobaltous  basic-carbonate,  Co805(C03)3 , 
peach-red,  which  when  boiled  loses  carbonic  anhydride  and  acquires  a 
violet,  or,  if  the  reagent  be  in  excess,  a  blue  color.  The  precipitate  is 
soluble  in  ammonium  carbonate  and  very  slightly  soluble  in  fixed  alkali 
carbonates.  Carbonates  of  Ba  ,  Sr ,  Ca  ,  or  Mg  do  not  precipitate  cobaltous 
chloride  or  nitrate  in  the  cold  (separation  from  Fe'",  Al,  and  Cr'"),  but 
by  prolonged  boiling  they  precipitate  them  completely.  However,  if  a 
solution  of  a  cobaltous  salt  be  treated  with  chlorine,  a  cobaltic  salt  is 
formed  (5a),  which  is  precipitated  in  the  cold  on  digestion  with  BaC03 
(distinction  from  Ni). 

6. — Oxalic  acid  and  oxalates  precipitate  reddish-white  cobaltous  oxalate, 
CoC2O4  ,  soluble  in  mineral  acids  and  in  ammonium  hydroxide. 

Alkali  cyanides — as  KCN — precipitate  the  brownish-white  cobaltous 
cyanide,  Co(CN)2 ,  soluble  in  hydrochloric  acid,  not  in  acetic  or  in  hydro- 
cyanic acid,  soluble  in  excess  of  the  reagent,  as  double  cyanides  of  cobalt 
and  alkali  metals — (KCN)2Co(CN)2 — potassium  cobaltous  cyanide,  the  solu- 
tion having  a  brown  color:  CoCl2  +  2KCN  =  Co(CN)2  +  2KC1 .  Then 
Co(CN)2  +  2KCN  =  (KCN),Co(CN)2 .  Dilute  acids,  without  digestion, 
reprecipitate  cobaltous  cyanide  from  this  solution  (the  same  as  with  Ni): 
(KCN)2Co(CN)2  +  2HC1  =  Co(CN)2  +  2HCN  +  2KC1 .  But  if  the  solu- 
tion, with  excess  of  the  alkali  cyanide  and  with  a  drop  or  two  of  Jiydro- 
chloric  acid,*  insuring  free  HCN ,  be  now  digested  hot  for  some  time,  the 
cobaltous  cyanide  is  oxidized  and  converted  into  alkali  cobalticyanide — as 
K3Co(CN)6 — corresponding  to  ferricyanides,  but  having  no  corresponding 
nickel  compound: 

4Co(CN)2  +  4HCN  +  O2  =  4Co(CN)3  (cobaltic  cyanide)  +  2H2O 
Co(CN)3  +  3KCN  =  K3Co(CN)6  (potassium  cobalticyanide). 

In  the  latter  solution  neither  alkalies  nor  acids  precipitate  the 
cobalt  (important  distinction  from  nickel,  whose  solution  remains 
(KCN)2Ni(CN)2 ,  and  which,  after  digestion  as  above  is  precipitated). 
The  potassium  cobalticyanide  solution,  after  removal  of  the  Ni ,  may  bo 
precipitated  with  HgN03  (Gibbs,  J.  C.,  1874,  27,  92).  The  oxidation 
of  the  cobalt  may  be  hastened  by  the  presence  of  chromic  acid,  whirl i  is 

*  Moore  (C.  N.,  1887,  56,  3)  adds  glacial  phosphoric  acid  to  the  neutral  solutions  of  cobalt  and 
nickel,  until  the  precipitate  first  formed  begins  to  redissolve ;  then  he  adds  KCN  and  boils, 
continuing  the  boiling- and  addition  of  KCN  until  K.OH  fails  to  give  a  precipitate.  He  then 
warms  with  excess  of  bromine  in  presence  of  KOH,  whereupon  the  nickel  is  completely  pre- 
cipitated leaving  the  cobalt  in  solution.  See  also  Hambly  (C.  N.<  1893,  65, 289). 


170  COBALT.  §132,  68. 

reduced  to  trivalent  chromium  compound:  6Co(CN)2  +  24KCN  + 
2Cr03  +  3H20  =  =  6K3Co(CN)6  +  Cr203  +  C>KOH  (McCulloch,  C.  N., 
1889,  59,  51),  also  by  means  of  NaCIO  or  NaBrO  produced  by  passing 
chlorine  into  caustic  soda  or  adding  bromide  to  the  alkali. 

Ferrocyanides,  as  K4Fe(CN)fi  ,  precipitate  colaltoits  fcn'ocyuiiirtc,  Co2Fe(CN)6  , 
gray-green,  insoluble  in  acids.  Ferricyanides,  as  K3Fe(CN)6  ,  precipitate  cobdU- 
(W#  ferricyanide,  Co3(Fe(CN)0)2  ,  brownish-red,  insoluble  in  acids.  But  a  more 
distinctive  test  is  made  by  adding-  ammonium  cMwide  and  lij/diwidc,  with  the 
ferricyanide,  when  a  blood-red  color  is  obtained,  in  evidence  of  cobalt  (distinc- 
tion from  nickel).  Potassium  xanthate  forms  a  green  precipitate  in  neutral  or 
slightly  acid  solutions  of  cobalt  salts  (§133,  66). 

Nitroso-^-naphthol  completely  precipitates  solutions  of  Cu ,  Fe ,  and  Co  ; 
Ag ,  Sn ,  and  Bi  salts  are  partially  precipitated;  and  Pb  ,  Hg ,  As  ,  Sb  ,  Cd  . 
Al ,  Cr ,  Mn ,  Ni ,  Zn ,  Ca ,  Mg ,  and  Gl  remain  in  solution  (Burgass,  Z. 
angew.y  1896,  596).  In  analysis  for  the  separation  of  cobalt  and  nickel  it  is 
recommended  to  proceed  as  follows :  The  solution  of  the  metals  preferably 
as  sulphates  or  chlorides  is  acidulated  with  hydrochloric  acid  and  treated 
with  a  hot  solution  of  nitroso-/3-naphthol  in  50  per  cent  acetic  acid,  until  the 
whole  of  the  cobalt  is  precipitated.  The  brick-red  precipitate  is  then  washed 
with  cold  HC1 ,  then  with  hot  12  per  cent  HC1 ,  and  finally  with  water.  The 
separation  is  quantitative.  The  precipitate  may  be  ignited  in  air  to  the 
oxide  or  with  oxalic  acid  in  an  atmosphere  of  hydrogen  and  weighed  as 
the  metal.  For  qualitative  purposes  the  cobalt  in  the  precipitate  may  be 
identified  by  the  color  of  the  borax  bead  (7).  The  nickel  in  the  filtrate 
may  be  precipitated  by  hydrosulphuric  acid  and  identified  by  the  usual 
tests  (Knorre,  B.,  1887,  20,  283  and  Z.  angeiv.,  1893,  264). 

Ammonium  TMocyanate  in  concentrated  solutions  of  cobalt  produces  a 
brilliant  blue  color  which  disappears  on  dilution  with  water.  On  the  ad- 
dition of  amyl  alcohol  and  ether  in  equal  portions  and  shaking,  the  layer 
of  alcohol  and  ether  becomes  blue.  The  test  is  very  delicate  especially  if 
the  cobalt  solution  is  concentrated.  Nickel  salts  produce  no  coloration  of 
the  amyl  alcohol  and  ether,  but  iron  interferes  because  the  red  Fe  (CNS)3 
is  dissolved  by  the  alcohol.  By  the  addition  of  2  or  3  c.c.  of  concentrated* 
ammonium  acetate  solution  and  2  or  3  drops  of  50  per  cent  tartaric  acid 
solution,  the  red  color  of  the  Fe  (CNS)3  may  be  removed.  The  blue  color 
is  probably  due  to  the  undissociated  salt  (NHJ2  [Co(CNS)4]  which  is 
soluble  in  ether  and  amyl  alcohol  (Vogel  Ber.  12,  2314.  Treadwell,  Z.  An. 
Hi.  26  (1901),  105). 

Ammonium  thiocyanate  may  be  employed  for  the  separation  of  nickel 
and  cobalt  (Rosenheim.  and  Cohen,  B.,  33,  111,  and  Rosenheim  and  Huld- 
shinsky,  B.f  34,  2050),  12  grams  of  NH4CNS  are  added  to  the  nitric  acid 
solution  of  the  metals  the  volume  being  15-20  c,c,  Cobalt  produces  a  deep 


§132,   6>.  COBALT.  171 

blue  compound  of  the  formula  R2Co(CNS)4  in  which  R  is  an  alkali  metal. 
This  blue  compound  may  be  extracted  by  a  mixture  of  25  vol.  ether  and 
1  vol.  amyl  alcohol. 

c. — Potassium  nitrite  forms  with  both  cobaltous  and  nickelous  salts  the 
double  nitrites,  Co(N02)2.2KNO,  and  Ni(N02),.2KN02 ,  soluble.  The  nickel 
compound  is  very  stable,  but  if  the  cobalt  compound,  strongly  acidulated 
with  acetic  acid,  be  warmed  and  allowed  to  stand  for  some  time,  preferably 
twenty-four  hours;  the  cobalt  is  completely  precipitated  as  the  yellow 
crystalline  potassium  cobaltic  nitrite,  Co(N02)3.3KH02  (separation  from 
Ni)  :  CoCl,  +  GKNO.,  +  HC2H302  +  HN02  =  Co(N02)3.3KN02  +  2KC1 
+  KC2H362  +  H20  +  NO. 

d. — Phosphates,  as  Na2HP04  ,  precipitate  cobaltous  salts  as  the  reddish 
cobalt  on  t>  i>Ito^i>Jiatc,  CoHPO4  ,  soluble  in  acids  and  in  ammonium  hydroxide. 
Sodium  pyrophosphate  forms  a  gelatinous  precipitate  with  solutions  of  cobalt 
salts,  soluble  in  excess  of  the  reagent.  The  addition  of  acetic  acid  causes  a 
precipitation  of  the  cobalt  even  in  the  presence  of  tartrates  (separation  from 
Ni,  but  not  from  Mn  or  Fe)  (Vortmann,  B.,  1888,  21,  1103).  If  a  solution  of 
cobaltous  salt  be  treated  with  a  saturated  solution  of  ammonium  phosphate 
and  hydrochloric  acid,  and  when  hot  treated  with  an  excess  of  ammonium 
hydroxide,  a  bluish  precipitate  of  CoNH4PO4  will  appear  on  stirring  (separa- 
tion from  nickel*)  (Clark,  C.  N.,  1883,  48,  262;  Hope,  J.  Soc.  Ind.,  1890,  9,  375). 

e. — Hydrosulphuric  acid,  with  normal  cobaltous  salts,  gradually  and 
imperfectly  precipitates  the  black  cobalt  sulphide,  CoS  ;  from  cobalt  acetate, 
the  precipitation  is  more  prompt,  and  is  complete;  but  in  presence  of 
mineral  acids,  as  in  the  second  group  precipitation,  no  precipitate  is  made. 
Immediate  precipitation  takes  place  with  hydrosulphuric  acid  acting  upon 
solutions  of  cobaltous  salts  in  ammonium  hydroxide.  When  formed,  the 
precipitate  is  scarcely  at  all  soluble  in  dilute  hydrochloric  acid  or  in  acetic 
acid;  slowly  soluble  in -moderately  concentrated  hydrochloric  acid;  readily 
soluble  in  nitric  acid;  and  most  easily  in  nitrohydrochloric  acid.  By 
exposure  to  the  air,  the  recent  cobaltous  sulphide  is  gradually  oxidized  to 
cobalt  sulphate,  soluble,  as  occurs  with  iron  sulphide  (§126,  60).  Alkali 
sulphides  precipitate  immediately  and  perfectly  the  black  cobaltous  sul- 
phide, described  above,  insoluble  in  excess  of  the  reagent.  When  cobaltous 
salts  are  boiled  with  sodium  thiosulphate  a  portion  of  the  cobalt  is  precipi- 
tated as  the  black  sulphide. 

f. — The  higher  oxides  of  cobalt  and  cobaltic  salts  are  reduced  by  warming 
with  halogen  acids,  liberating  the  corresponding  halogens  (HC1  does  not  reduce 
the  cobalt  in  K,Co(CN)8). 

g. — Soluble  arsenites  and  arsenates  precipitate  cobaltous  salts,  forming  the 
corresponding  cobaJt  iirxenites  or  arsenates,  bluish-white,  soluble  in  ammonium 
hydroxide  or  in  acids,  including1  arsenic  acid.  //. — Soluble  chromates  precipi- 
tate colmltous  eliminate,  yellowish-brown,  soluble  in  ammonium  hydroxide  and 

*  Krauss  (Z.,  1891,  3O,  227)  gives  a  good  review  of  the  most  important  methods  for  the  separa- 
tion of  cobalt  and  nickel. 


COBALT.  §132,  6/. 

in  acids,  including  chromic  acid.  No  precipitate  is  formed  with  potassium 
dichromate.  i. — KMnO4  added  to  an  annnoniaeal  solution  of  eobaltous  salts 
oxidizes  the  cobalt  and  prevents  its  precipitation  by  KOH  (separation  from 
Ni)  (Delvaux,  G.  r.,  1881,  92,  723). 

;'. — Cobaltous  salts  in  ammoniacal  solution,  warmed  with  H2O2  and  then 
rendered  acid  with  acetic  acid,  are  precipitated  by  ammonium  molybdate 
(separation  from  Ni)  (Carnot,  C.  r.,  1889,  109,  109). 

7.  Ignition, — In  the  bead  of  borax.,  and  in  that  of  microcosmic  salt,  with 
oxidizing  and  with  reducing  flames,  cobalt  gives  an  intense  blue  color. 
The  blue  bead  of  copper  changes  to  brown  in  the  reducing  flame.     If 
strongly  saturated,  the  bead  may  appear  black  from  intensity  of  color,  but 
will  give  a  blue  powder.     This  important  test  is*  most  delicate  with  the 
borax  bead.     Manganese,  copper,  nickel,  or  iron  interfere  somewhat.     By 
ignition,  with  sodium  carbonate  on  charcoal  or  with  the  reducing  flame, 
compounds   of  cobalt  are  reduced  to  the  metal  (magnetic).     Cobaltous 
oxide  dissolves  in  melted  glass  and  in  other  vitreous  substances,  coloring 
the  mass  blue — used  to  cut  off  the  light  of  yellow  flames  (§205,  7).     The 
black  cobaltoso-cobaltic  oxide,  Co304 ,  as  left  by  ignition  of  eobaltous  oxide 
or  nitrate,  combines  or  mixes,  by  ignition,  with  zinc  oxide  from  zinc  com- 
pounds to  form  a  green  mass,  with  aluminum  compounds  a  blue,  and  with 
magnesium  compounds  a  pink  mass. 

8.  Detection. — After  removal  of  the  metals  of  the  first  three  groups 
cobalt  is  precipitated  by  H2S  in  ammoniacal  solution  with  Ni ,  Mn  and  Zn  . 
The  sulphides  are  digested  with  cold  dilute  HC1  which  dissolves  the  Mn 
and  Zn  .     The  borax  bead  test  (7)  is  now  made  upon  the  remaining  black 
precipitate,  and  if  Ni  be  not  present  in  great  excess  *  the  characteristic  blue 
bead  is  obtained.    If  the  nickel  be  present  in  such  quantities  as  to  obscure 
the  blue  borax  bead  the  sulphides  are  dissolved  in  hot  cone.  HC1,  using  a 
few  drops  of  HN03  .     The  solution  is  heated  to  decompose  all  the  nitric 
acid  and,  after  dilution,  the  cobalt  is  precipitated  with  nitroso-/?-naphthol, 
according  to  directions  given  in  6&,  and  further  identified  by  the  bead  test. 

The  NH4CNS  test  (§132,  6&)  and  the  NaHC03  and  H202  test  (§140) 
may  also  be  obtained. 

9.  Estimation. — (1)  As  metallic  cobalt,  all  compounds  that  may  be  reduced 
by  ignition  in  hydrogen  gas,    e.  g.  CoCl2 ,   Co(NO3)2,  CoCO3 ,  and  all  oxides  and 
hydroxides.      (2}  As  UoD  ,  all  soluble  cobalt  salts,  all  salts  whose  acids  are  expelled 
or  destroyed  by  ignition,  all  oxides  and  hydroxides.     The  salt  is  converted  into 
Co(OH)2  by  precipitation  with  a  fixed  alkali,  and  ignited  in   a    stream    of  CO2 . 
The  carbonate  and  nitrate  may  be  ignited  directly    in    CO2,  and   organic    salts 
are  first  igni  ed  in  the  air  until  the   carbon  is  oxidized,   and  then  again  ignited 
in  CO2 .      (3)   After    converting   into  a  sulphate  it  is  ignited  at  a  dull-red  heat 
and  weighed   as    a    sulphate.      (4)  After    converting    into    the   oxalate,  titrated 
with    KMnCV      (5)  In  presence  of  nickel,  it  is  oxidized  in  alkaline  solution  by 
H2O2,  KI  and  HCl  are  added,    and    the    liberated    iodine  titrated  with  sodium 
thiosulphate    (Fischer,  C.   C.,    1889,   116).      (6)  Electrolytically.       (7)  Separated 

*  If  more  than  ten  parts  of  nickel  are  present  to  one  part  of  cobalt,  the  characteristic  blue 
bead  is  not  obtained. 


§133,  5b.  NICKEL.  173 

from  nickel  by  nitroso-/3-naphthol,  and  after   ignition  in  hydrogen  weighed  as 
the  metal  (66). 


10.  Oxidation. — Co"  is  oxidized  to  Co'"  in  presence  of  a  fixed  alkali  by 
Pb02 ,  Cl ,  KC10 ,  Br ,  KBrO ,  I  and  H202* ;  in  presence  of  acetic  acid  by 
KN02  (Gc).  Co'"  is  reduced  to  Co"  by  H2C204 ,  H,P02 ,  H2S  ,  H2S03 ,  HC1 , 
HBr ,  and  HI .  Metallic  cobalt  is  precipitated  from  solution  of  CoCl2  by 
Zn  ,  Cd  .  and  Mg  . 

§133.  Niekel.     Ni  =  58.70  .     Usual  valence  two  and  three. 

1.  Properties.— Specific  gravity,  8.9  (Rchroeder,  Pogff.,  1850,  106,  226).     Melting 
point,  1452°  (Cir.  B.  S.,  36,  1915).     It  is  a  hard  white  metal,  capable  of  taking 
a  high  polish;    malleable,  ductile  and  very  tenacious,  forming  wire  stronger  than 
iron  but  not  quite  so  strong  as  cobalt  (§  132,   1).     It  does  not  oxidize  in  dry  or 
moist   air   at   ordinary   temperatures.     It   is   magnetic   but   loses   its   magnetism 
like  steel  on  heating  to  redness  (Gangain,  C.  r.,  1876,  83,  661).     It  burns  with 
incandescence  when  heated  in  O ,  Cl ,  Br ,  or  S .     It  is  much  used  in  plating  other 
metals,  in  making  coins  of  small  denominations,  in  hardening  armor  plate,  pro- 
jectiles, etc.     The  presence  of  small  amounts  of  phosphorus  or  arsenic  renders  it 
much   more   fusible,    without   destroying   its   ductility;     a   larger   amount   makes 
it  brittle. 

2.  Occurrence. — Nickel  almost  always  occurs  in  nature  together  with  cobalt. 
It  is  found  as  millerite,  (NiS)  ;  pentlandite  (FeNi)S  ;    niccolite  (NiAs)  ;  garnierite 
(variable,  perhaps  H2(NiMg)SiO4  +  xH2O)  ;     frequently  in  pyrrhotite,   (FenSn+i 
with  Ni)  and  in  numerous  other  rarer  minerals. 

3.  Preparation.—  (1)  By  electrolysis.     (2)  By  heating-  jn  a  stream  of  hydrogen. 
The  oxide  is  reduced  in  this  manner  a't  270°   (W.  Miiller,  Pogg.,  1869,  136,  51). 

(3)  By    fusing-    the    oxalate    under    powdered    glass    (C02    being    given    off). 

(4)  Reduction  by  igniting  in  CO  .     (5)   Reduction  by  fusing  with  carbon  in  a 
variety  of  methods.     (6)  By  heating  the  carbonyl,  f  Ni(CO)4  to  200°. 

4.  Oxides  and  Hydroxides. — Nickelous  oxide  is  formed  when  the  carbonate, 
nitrate,   or   any   of   its   oxides   or   hydroxides   are   strongly   ignited.     Nickelous 
hydroxide   is    formed    by    precipitation    of    nickelous    salts    with   fixed    alkalis. 
Nickclic  oxide,  Ni203  ,  is  made  from  NiC03  ,  Ni(NO3)2  or  NiO  by  heating  in  the 
air  not  quite  to  redness,  with  constant  stirring.     It  is  changed  to  NiO  at  a  red 
heat.     Nickelic   hydroxide,    Ni(OH)3  ,    is   formed    by   treating   nickelous    salts 
first  with  a  fixed  alkali  hydroxide  or  carbonate  and  then  with  Cl ,  NaCIO  ,  Br 
or  NaBrO  (not  formed  by  iodine),  a  black  powder  forming  no  corresponding 
salts  (Campbell  and  Trowbridge,  J.  Anal.,  1893,  7,  301).     A  trinickelic  tetroxide, 
Ni3O4  ,  magnetic  (corresponding  to  Co3O4  ,  Fe3O4  ,  Mn3O4  and  Pb3O4),  is  formed, 
according  to  Baubigny  (C.  r.,  1878,  87,  1082),  by  heating  NiCL  in  oxygen  gas 
at  from  350°  to  440°;  and  by  heating  Ni2O3  in  hydrogen  at  190°  (Moissan,  A.  Ch., 
1880,  (5),  21,  199). 

5.  Solubilities. — a. — Metal. — Hydrochloric    or    sulphuric    acid,    dilute    or    con- 
centrated, attacks  nickel  but  slowly  (Tissier,  C.  r.,  1860,  50,  106);  dilute  nitric 
acid  dissolves  it  readily,  while  towards  concentrated  nitric  acid  it   acts  very 
similar  to  passive  iron  (Deville,  C.  r.,  1854,  38,  284).     It  is  not  attacked  when 
heated  in   contact  with   the   alkali   hydroxides  or  carbonates,     ft. — Oxides  and 

*  Durant,  C.  N.,  1897,  75,  43. 

t  Nickel  carbonyl  is  prepared  by  heating  the  nickel  ore  in  a  current  of  CO  .  It  is  a  liquid,  sp. 
gr.  1.3185,  boiling  at  43°  and  freezing  at  —25°.  When  heated  to  200°  it  is  decomposed  into  Ni 
#nd  CO  (Berthelot,  C,  r,,  1891,  112,  1343;  113,  679;  Mond,  J.  Soc,  Ind.,  1892.  11,  750). 


174  NICKEL.  §133,  firt. 

hydroxides. —  Nickelous  oxide  and  hydroxide  are  insoluble  in  water  or  fixed 
alkalis,  soluble  in  ammonium  hydroxide  and  in  acids.  Nickelie  oxides  and 
hydroxides  are  dissolved  by  acids  with  reduction  to  nickelous  salts,  with  halogen 
acids  the  corresponding-  halogens  are  liberated.  The  moist  nickelic  hydroxide, 
formed  by  the  action  of  Cl  ,  Br  ,  etc.,  in  alkaline  solution,  after  washing  with 
hot  water  liberates  free  iodine  from  potassium  iodide  (distinction  from  cobalt). 
Nickelic  hydroxide  when  treated  with  dilute  sulphuric  acid  forms  NiSO4  , 
oxygen  being  evolved.  With  nitric  acid  the  action  is  similar,  distinction  from 
cobaltic  hydroxide,  which  requires  a  more  concentrated  acid  to  effect  a  similar 
reduction,  c. — Salts. — The  salts  of  nickel  have  a  delicate  green  color  in  crystals 
and  in  solution;  when  anhydrous,  they  are  yellow.  The  nitrate  and  chloride 
are  deliquescent  or  efflorescent,  according  to  the  hygrometric  state  of  the 
atmosphere;  the  acetate  is  efflorescent.  The  chloride  vaporizes  at  high  tem- 
peratures. 

The  carbonate,  sulphide,  phosphate,  borate,  oxalate,  cyanide,  ferrocyanide 
and  ferricyanide  are  insoluble;  the  double  cyanides  of  nickel  and  alkali 
metals,  soluble  in  water.  The  chloride  is  soluble  in  alcohol,  and  the  nitrate  in 
dilute  alcohol.  Most  salts  of  nickel  form  soluble  compounds  by  action  of 
ammonium  hydroxide. 

6.  Reactions,  a. — Alkali  hydroxides  precipitate  solutions  of  nickel 
salts  as  nickel  hydroxide,  Ni(OH)2 ,  pale  green,  not  oxidized  by  exposure  to 
the  air  (§132,  6a),  insoluble  in  excess  of  the  fixed  alkalis  (distinction  from 
zinc),  soluble  in  ammonium  hydroxide  or  ammonium  salts,  forming  a 
greenish-blue  to  violet-blue  solution.  Excess  of  fixed  alkali  hydroxide 
will  slowly  precipitate  nickel  hydroxide  from  the  ammoniacal  solutions 
(distinction  from  cobalt).  Alkali  carbonates  precipitate  green  'basic 
nickelous  carbonate,  Ni5(OH)6(C03)2  (composition  not  constant),  soluble  in 
ammonium  hydroxide  or  ammonium  salts,  with  blue  or  greenish-blue  color. 
Carbonates  of  Ba ,  Sr ,  Ca ,  and  Mg  are  without  action  on  nickelous 
chloride  or  nitrate  in  the  cold  (distinction  from  Fe"',  Al ,  and  Cr'"),  but 
on  boiling  precipitate  the  whole  of  the  nickel. 

ft. — Oxalic  acid  and  oxalates  precipitate,  very  slowly  but  almost  completely, 
after  twenty-four  hours,  nickel  oxalate,  green.  Alkali  cyanides,  as  KCN  ,  pre- 
cipitate nickel  cyanide,  Ni(CN)2  ,  yellowish-green,  insoluble  in  hydrocyanic 
acid,  and  in  cold  dilute  hydrochloric  acid;  dissolving  in  excess  of  the  cyanide, 
by  formation  of  soluble  double  cyanides,  as  potassium  nickel  cyanide 
(KCN)2Ni(CN)2  .  The  equation  of  the  change  corresponds  exactly  to  that  for 
cobalt  (§132,  66);  and  the  solution  of  double  cyanide  is  reprecipitated  as 
Ni(CN)2  by  a  careful  addition  of  acids  (like  cobalt);  but  hot  digestion,  with 
the  liberated  hydrocyanic  acid,  forms  no  compound  corresponding  to  cobalti- 
cyanides,  and  does  not  prevent  precipitation  by  acids  (distinction  from  cobalt). 
It  will  be  observed  that  excess  of  hydrochloric  or  sulphuric  acid  will  dissolve 
;he  precipitate  of  Ni(CN)2  .  If  an  oxidizing  agent  such  as  NaCIO  or  NaBrO 
be  added  to  the  alkaline  solution  of  the  double  cyanide,  the  nickel  will  be  oxidized 
and  precipitated  (separation  from  cobalt),  as  Ni2O3.3H2O ,  according  to  the 
following  equation: 

K2Ni(CN)4  +  NaBrO  +  NaOH  =  Ni2O3  3H2O 

Ferrocyanides,  as  K4Fe(CN)6 ,  precipitate  a  greenish-white  nickel  ferrocyanide, 
Ni-Fe(CN)6  ,  insoluble  in  acids,  soluble  in  ammonium  hydroxide,  decomposed 
by  fixed  alkalis.  Ferricyanides  precipitate  greenish-yellow  nickel  ferricyanide, 
insoluble  in  acids,  soluble  in  ammonium  hydroxide  to  a  green  solution  (§132.  6b), 


§133,  7.  NICKEL.  175 

A  solution  of  nitre  ferricyanide  precipitates  solutions  of  cobalt  and  nickel  salts, 
the  latter  being  soluble  in  dilute  ammonium  hydroxide  (Cavalli,  Gazzetta,  1897, 
27,  ii,  95). 

A  solution  of  potassium  xanthate  precipitates  neutral  solutions  of  nickel  and 
cobalt,  the  former  being-  soluble  in  ammonium  hydroxide  (distinction),  from 
which  solution  it  is  precipitated  by  (NH4),S  (Phipson,  C.  N.,  1877,  36,  150). 
The  xanthate  also  precipitates  nickel  in  alkaline  solution  in  presence  of 
Na4P,0T  (a  separation  from  Fe'")  (Campbell  and  Andrews,  /.  Am.  Soc.,  1895, 
17,  125). 

Xirl-rl  salts  are  not  precipitated  by  an  acetic  acid  solution  of  nitroso-/?- 
naphthol  (separation  from  cobalt)  (Knorre,  B.,  1885,  18,  702). 

C-  —  Potassium  nitrite  in  presence  of  acetic  acid  does  not  oxidize  nickelous 
compounds  (distinction  from  cobalt),  d.  —  Sodium  phosphate,  Na2HPO4  ,  pre- 
cipitates nickel  phosphate,  Ni.,(P04)2  ,  greenish-white. 

0-  —  Hydrosulphuric  acid  precipitates  from  neutral  solutions  of  nickel 
salts  a  portion  of  the  nickel  as  nickel  sulphide,  black  (Baubigny,  C.r..  1882, 
94,  1183;  95,  34).  The  precipitation  takes  place  slowly,  and  from  nickel- 
ous acetate  is  complete.  In  the  presence  of  mineral  acids  no  precipita- 
tion takes  place.  Alkali  sulphides  precipitate  the  whole  of  the  nickel, 
as  the  black  sulphide.  Although  precipitation  is  prevented  by  free  acids, 
the  precipitate,  once  formed,  is  nearly  insoluble  in  acetic  or  in  dilute 
hydrochloric  acids;  slowly  dissolved  by  concentrated  hydrochloric  acid, 
readily  hy  nitric  or  nitro-hydrochloric. 

Nickel  sulphide,  NiS  ,  is  partially  soluble  in  yellow  ammonium  sulphide,* 
from  which  brown-colored  solution  it  is  precipitated  (gray,  black  mixed  with 
sulphur)  on  addition  of  acetic  acid  (distinction  from  cobalt).  It  is  insoluble 
in  the  monosulphide  and  is  completely  precipitated  by  passing  hydrogen  sulphide 
through  an  ammoniacal  solution  in  the  absence  of  air  or  oxidizing  agents  (Noyes, 
Bray  and  Spear,  J.  Am.  Soc.,  30,  497).  Freshly  precipitated  nickel  sulphide 
is  soluble  in  KCN  and  reprecipitated  as  Ni(CN)2  on  adding  HC1  or  H2SO4  (sep- 
aration from  cobalt)  (Guyard,  Bl.,  1876,  (2),  25,  509).  When  nickel  salts  are 
boiled  with  a  solution  of  NfeSzOs,  a  portion  of  the  nickel  is  precipitated  as  the 
black  sulphide. 

f-  —  The  halogen  acids  reduce  the  higher  oxides  of  nickel  to  nickelous 
salts  with  liberation  of  the  corresponding  halogen.  Potassium  iodide 
added  to  freshly  precipitated  nickelic  hydroxide  gives  free  iodine  (distinc- 
tion from  cobalt). 

9-  —  Nickel  salts  are  precipitated  by  arsenites  and  arsenates,  white  or  green- 
ish-white, soluble  in  acids,  including  arsenic  acid.  h.  —  Potassium  chromate 
precipitates  basic  nickel  chromate,  yellow,  soluble  in  acids,  including  chromic 
acid  (Schmidt,  A.,  1870,  156,  19).  K,Cr2OT  forms  no  precipitate. 

7.  Ignition.  —  Xickel  compounds  dissolve  clear  in  the  borax  bead,  giving1  with 
the  oxidizing  flame  a  purple-red  or  violet  color  while  hot,  becoming  yellowish- 
brown  when  cold;  with  the  reducing1  flame,  fading-  to  a  turbid  gray,  from 
reduced  metallic  nickel,  and  finally  becoming-  colorless.  The  addition  of  any 
potassium  salt,  as  potassium  nitrate,  causes  the  borax  bead  to  take  a  dark 
purple  or  blue  color,  clearest  in  the  oxidizing-  flame.  With  microcosmic  salt, 

*  Hare  (J".  Am.  Soc.,  1895,  17,  537)  adds  tartaric  acid  to  the  solutions  of  nickel  and  cobalt,  and  an 
excess  of  sodium  hydroxide.  He  then  passes  in  IF2S.  The  cobalt  is  completely  precipitated 
while  the  nickel  remains  in  solution,  and  can  be  precipitated  upon  acidulating-  the  nitrate. 


ft 


f*HI  Lf-BK 


1?6  X1CK&L.  §133,  8. 

nickel  gives  a  reddish-brown  bead,  cooling  to  a  pale  reddish-yellow,  the  rob- 
being  alike  in  both  flames.  Hence,  with  this  reagent,  in  the  reducing  flame, 
the  color  of  nickel  may  be  recognized  in  presence  of  iron  and  manganese,  which 
are  colorless  in  the  reducing  flame;  but  colxilt  effectually  obscures  the  bead 
test  for  nickel.  The  yellow-red  of  copper  in  the  reducing  flame,  persisting  in 
beads  of  microcosmic  salt,  also  masks  the  bead  test  for  nickel.  By  ignition 
with  sodium  carbonate  on  charcoal,  compounds  of  nickel  are  reduced  to  the 
metal,  slightly  attracted  by  the  magnet. 

8.  Detection. — We  proceed  exactly  as  with  cobalt  for  the   nitroso-/3 
naphthol  precipitation.     The  Ni  remains  in  the  filtrate  and  can  be  precipi 
tated  with,  H2S  (after  neutralizing  with  NH4OH),  and  its  presence  con- 
firmed by  the  usual  tests.     Or  dissolve  the  sulphides  of  Ni  and  Co  io 
HN03 ,  evaporate  nearly  to  dryness,  add  an  excess  of  KOH  or  Na2C03 , 
boil,  add  bromine  water  and  boil  to  complete  oxidation  of  the  Co  and  Ni  , 
filter,  wash  thoroughly  with  hot  water  and  add  hot  solution  of  KI  to  the 
precipitate  on  the  filter  paper.     Free  iodine  (test  with  CS2)  is  evidence  of 
the  presence  of  nickel. 

Nickel  may  also  be  detected  in  the  presence  of  cobalt  as  follows : — 
Dissolve  the  sulphides  of  Ni  and  Co  in  aqua  regia.  boll  out  the  excess  of 
chlorine,  neutralizing  with  KOH  and  add  KCN  in  slight  excess.  Add 
NaCIO  or  NaBrO  and  warm.  The  nickel  is  percipitated  as  the  brown 
hvdrated  oxide  Ni,03.3H.,0  . 

Dimethylglyoxime  produces  a  red  crystalline  precipitate  which  forms 
the  most  sensitive  test  known  for  detecting  nickel  in  the  presence  of  cobalt. 
The  reagent  is  prepared  by  dissolving  1  gram  of  the  dimethyglyoxime  in 
100  c.c.  of  98%  alcohol.  The  solution  should  be  made  slightly  alkaline 
with  ammonia  and  boiled  after  adding  a  few  drops  of  the  reagent.  The 
following  reactions  take  place : 

2(CH3)2C2N2O2H2+NiCl2  -f  2NH4OH  =  2NH4C1  +  [(CH3)2C2N2O2H]2Ni  +  2H2O 

If  the  amount  of  nickel  is  small  the  solution  at  first  becomes  yellow 
and  on  cooling  deposits  red  needles.  The  test  is  sensitive  to  nickel  when 
present  in  400,000  parts  of  water  (L.  Tschugaeff  Ber.  38  (1905),  2520). 
Ten  times  as  much  cobalt  may  be  present  but  in  the  presence  of  larger 
quantities  of  cobalt  the  following  procedure  is  followed.  Excess  of  am- 
monia is  added  to  the  cobalt  solution  and  then  a  few  cubic  centimeters 
of  hydrogen  peroxide.  The  solution  is  boiled  to  decompose  the  excess  of 
hydrogen  peroxide.  Dime  thy  Iglyoxime  is  added  and  the  solution  again 
boiled.  If  a  small  amount  of  nickel  is  present  a  red  scum  is  formed  on 
the  section  and  red  crystals  form  on  the  sides  of  the  beaker.  On  filtering 
the  solution  the  red  precipitate  L;  readily  observed  on  the  filter  paper. 

9.  Estimation. * — (1}  Nickel  hydroxide,  oxide,  carbonate  or  nitrate  is  ignited 
at  a  white  heat  and  weighed  as  NiO .     (2}  It  is  converted  into  the  sulphate  and 

*  Goulal  (Z.  angew.,  1898,  177)  gives  a  summary  of  the  methods  proposed  for  the  volu- 
metric estimation  of  nickel. 


§134,  4.  MANGANESE.  177 

deposited  on  platinum  as  the  free  metal  by  the  electric  current.  (-?)  Volu- 
metrically.  By  lilration  In  a  slightly  alkaline  solution  with  KCN  ,  using-  a 
small  amount  of  freshly  precipitated  Agl  as  an  indicator  (Campbell  and 
Andrews,  J.  Am.  ftoc.,  1895,  17,  127). 

10.  Oxidation. — Ni"  is  changed  to  Hi'"  in  presence  of  fixed  alkalis  b} 
Cl ,  NaCIO  ,  Br ,  and  NaBrO  (not  by  I ,  distinction  from  cobalt,  Donatti, 
B.,  1879,  12,  1868).  Ni'"  is  reduced  to  Ni"  by  all  non-reducing  acids  with 
evolution  of  oxygen;  by  reducing  acids,  H,C,,04  is  oxidized  to  C02 ,  HNO? 
to  HN03 ,  H3P02  to  H3P04 ,  H2S  to  S ,  H,S03  to  H2S04 ,  HC1  to  Cl ,  HBr  to 
Br ,  HI  to  I ,  HCNS  to  HCN  and  H2S04 ,  H4Fe(CN)6  to  H3Fe(CN)6 .  Ni" 
is  reduced  to  the  metal  by  finely  divided  Zn  ,  Cd  ,  and  Sn  . 


§134.  Manganese.     Mn  =  54.93.    Valence    two,    three,    four,    six    and 

seven. 

1.  Properties. — Specific    gravity,    7.392    (Glatzel,    Ber.,   1889).     Melting    point, 
1260°  (C'ir.  B.  S.,  35,  1915).     Boiling  point,  1900°  (v.  d.  Weyde,  Ber.  44,   1879). 
It  is  a  brittle  metal,  having  the  general  appearance  of  cast  iron,  non-magnetic, 
takes  a  high  polish.     According  to  Deville  it  has  a  reddish  appearance.     It  is 
readily  oxidized,  decomposing  water  at  but  little  above  the  ordinary  temperature 
(Deville,   A.  Ch.,   1856,  (3),   46,   199).     It   is   used  largely   as  ferromanganese  in 
the   manufacture    of    Bessemer    steel. 

Oxides  and  hydroxides  of  manganese  exist  as  dyad,  triad,  and  tetrad;  the 
salts  exist  most  commonly  as  the  dyad  with  some  unstable  triad  and  tetrad 
salts;  as  an  acid  it  is  a  hexad  in  manganates  and  a  heptad  in  permanganates. 

2.  Occurrence. — Not  found  native.     It  accompanies  nearly  HI  iron  ores.     Its 
chief  ore  is  pyrolusite  (MnO-2)  .     It  is  Hso  found  as  braunite,  (8Mn>O3.MnSiO;i)  ; 
hausmannite,   (Mn;iO4)  ;    manganite,  (Mn2O3.H2O)  ;    rhodocrosite,  MnCO;i  ;    ala- 
bandite,  (MnS)  ;   and  as  a  constituent  of  many  other  minerals. 

3.  Preparation. — (/)  By  electrolysis  of  the  chloride.     (2)  By  reduction  with 
metallic  sodium  or  magnesium  (Glatzel,  B.,  1889,  22,  2857).     (3)   By  reduction 
with   some  form  of  carbon.     It  has  not  been   reduced    by   hydrogen.     (4)    By 
ignition  with  aluminum  (Goldschmidt,  A.,  1898,  301,  19). 

4.  Oxides  and  Hydroxides. — (a)  Mangunous  oxide,  MnO  ,  represents  the  only 
base  capable  of  forming  stable  manganese  salts.     It  is  formed   (1)   by  simple 
ignition  of  Mn(OH)2  ,  MnC03  or  MnC,04  ,  air  being  excluded;   (2)  by  ignition 
of  any  of  the  higher  oxides  of  manganese  with  hydrogen  in  a  closed   tube 
(Moissan,  A.  Ch.,  1880,   (5),  21,  199).     If  prepared  at  as  low  a  temperature  as 
practicable,  it  is  a  dark  gray  or  greenish-gray  powder,  and  oxidizes  quickly 
in  the  air  to  Mn304  .     If  prepared  at  a  higher  heat  it  is  more  stable.     Man- 
ganous  hydroxide,  Mn(OH)2  ,  is  formed  from  manganous  salts  by  precipita- 
tion  with   alkalis."  It   quickly   oxidizes   in   the    air,    forming   MnO  (OH),    thus 
changing  from  white  to  brown.     (?>)   Manganic  oxide,   Mn2O3  ,  is  formed  by 
heating  any  of  the  oxides  or  hydroxides  to  a  red  heat  in  oxygen  gas  or  in  air 
(Schnieder,  Pogg.,  1859,   107,  605).     Mang-anic  oxide-hydroxide,  MnO  (OH)  ,  is 
formed   (1)  by  oxidation  of  Mn(OH)2  in  the  air;   (2)  by  treating  MnO,  with 
concentrated  H,SO4  at  a  temperature  of  about  130°,  forming  Mn2(S04)3   and 
then  adding  water:  Mn2(SO4)3   +  4H20  =  2MnO(OH)    +  3H2SO4    (Carius,  A., 
1856,  98,  63).     (c)  TrimanganesQ  tetroxide,  Mn3O4  ,  is  formed  when  any  of  the 
higher  or  lower  oxides  of  manganese  or  any  manganese  salts  with  a  volatile 
acid  are  heated  in  the  air  to  a  white  heat  (Wright  and  Luff,  B.,  1878,  11,  2145). 
The  corresponding  hydroxide  would  be  Mn3(OH)8;  this  has  not  been  isolated. 
A   corresponding   oxide-hydroxide   is   formed    by    adding   freshly    formed    and 
moist  Mn02  to  an  excess  of  MnCl2  containing  NH4C1  (Otto,  A.,  1855,  93,  372). 


178  MANGANESE.  §134,  5. 

(W)  Manganese  peroxide,  MnO2 ,  is  formed  (/)  by  heating  Mn(NO.02  to  200° 
(Gorgeu,  C.  r.,  1879,  88,  796):  (2)  by  heating1  MnCO.,  with  KC103  to  300°;  (5)  by 
boiling-  any  manganous  salt  with  concentrated  HNO3  and  KC1O3 .  A  correspond- 
ing- hydroxide,  Mn(OH)4  ,  has  not  been  isolated.  Several  other  hydroxides, 
e.  ff.,  MnO(OH)2  ,  Mn,O3(OH)2  ,  Mn:,04(OH)4  etc.,  have  been  produced.  The 
chief  use  of  mang-anese  dioxide  is  in  the  preparation  of  chlorine  or  bromine. 
(e)  TVEang-anates. — Manr/anic  acid,  H2MnO4  ,  is  not  known  in  a  free  state.  The 
corresponding  salt.  K2Mn04  .  is  formed  when  any  form  of  manganese  is  fused 
with  KOH  or  K,CO3  (1)  in  the  air,  oxygen  being  absorbed:  or  (2)  with  KN03 
or  KC1O3  ,  NO  or  KC1  being  formed.  A  manganate  of  the  alkali  metals  is 
soluble  in  water,  with  gradual  decomposition  into  manganese  dioxide  and  per- 
manganates: ?,K2Mn04  +  2H,O  =  2KMnO.  -f  Mn02  +  4KOH  .  Free  alkali 
retards,  and  free  acids  and  boiling  promote,  .this  ehang-e.  Manganates  have 
a  green  color,  which  turns  to  the  red  of  permanganates  during  the  decomposi- 
tion inevitable  in  solution.  This  is  the  usual  method  of  manufacturing  KMn04 . 
(/)  Permanganic  acid  is  not  in  use  as  an  acid,  but  is  represented  by  .the  per- 
mang-anates,  as  KMnO4  .  The  permanganic  acid  radical  is  at  once  decomposed 
by  addition  of  hot  H2SO4  to  a  solid  permanganate  (1),  but  in  water  solution 
this  decomposition  does  not  at  once  take  place,  except  by  contact  with  oxidiz- 
able'  substances.  The  oxidizing  power  of  permanganates  extends  to  a  great 
number  of  substances,  possesses  different  characteristics  in  acid  and  in  alka- 
line solutions,  and  acts  in  many  cases  so  rapidly  as*  to  be  violently  explosive. 
The  reactions  with  ferrous  salts  (2)  and  with  oxalic  acid  (3)  are  much  used  in 
volumetric  analysis. 

(1)  4KMnO4  +  2H2SO4  —  2K2S04  +  4Mn02  +  302  +  2H20 
and  2Mn02  -f  2H2S04  =  2MnS04  +  2H20  +  02 

or  4KMn04  +  6H2S04  =  4MnS04  +  2K2SO4  +  502  +  GH20 

(2)  KMn04  +  SFeCL  +  SBttl  =  MnCL  +  KC1  +  5FeCl3  +  4H2O 

(3)  2KMn04  -f  5H2C2O4  +  6HC1  =  2MnCl2  +  2KC1  +  SH20  -f  10C02 

5.  Solubilities. — ff. — Metal. — Manganese  dissolves  readily  in  dilute  acids  to 
form  manganous-salts.  Concentrated  H2S04  dissolves  it  only  on  warming,  302 
being1  evolved.  It  combines  readily  with  chlorine  and  bromine,  ft. — Oxides 
and  hydroxides. — All  oxides  and  hydroxides  of  manganese  are  insoluble  in 
water.  They  are  soluble,  upon  warming,  in  hydrochloric  acid,  forming1  man- 
ganous  chloride;  the  higher  oxides  and  hydroxides  being  reduced  with  evolu- 
tion of  chlorine  (commercial  method  of  preparation  of  chlorine).  Instead  of 
hydrochloric  acid,  sulphuric  ac'd  and  a  chloride  may  be  employed  (HBr  and 
HI  act  similarly  to,  and  more  readily  than  HC1).  In  the  cold,  hydrochloric 
acid  dissolves  MnO2  to  a  greenish-brown  solution,  containing,  probably,  MnCl3 
or  MiiCl4  ,  unstable,  giving  chlorine  when  warmed  and  forming  Mn02  when 
strongly  diluted  with  water  (Pickering,  J.  C.,  1879,  35,  654;  Nickles,  A.  C7i-., 
1865,  (4),  5,  161).  Nitric  and  sulphuric  acids  dissolve  manganous  oxide  and 
hydroxide  to  mangjanous  salts.  Manganese  dioxide  (or  hydrated  oxide)  is 
insoluble  in  nitric  acid,  dilute  or  concentrated;  concentrated  sulphuric  acid 
with  heat  decomposes  it,  evolving-  oxygen  and  forming  manganous  sulphate: 
2MnO2  +  2H2SO4  =  2MnSO4  +  2H2O  +  O2  .  Manganous  hydroxide  is  insoluble 
in  the  alkalis  but  soluble  m  solutions  of  ammonium  salts. 

c. — Salts. — Manganous  sulphide,  carbonate,  phosphate,  oxalate,  borate, 
and  sulphite  are  insoluble  in  water,  readity  soluble  in  dilute  acids.  Man- 
ganic salts  are  somewhat  unstable  compounds,  of  a  reddish-brown  or 
purple-red  color,  becoming  paler  and  of  lighter  tint  on  reduction  to  the 
manganous  Combination.  MnCL  and  MnS04  arc  (1clii~/uc$cent.  Man- 
ganic chloride,  MnCl3,  and  the  percliloridc  ,  MnCl4,  are  unstable  salts  which 


§134,  6fl.  MANGANESE.  179 

are  decomposed  by  water  especially  when  hot,  to  MnCL  and  chlorine. 
The  trichloride  is  greenish  black,  soluble  in  absolute  alcohol  and  ether; 
while  the  tetrachloride  is  reddish-brown  soluble  in  absolute  alcohol.  (W. 
I>.  Holmes,  J.  Am.  tioc.  29,  1277.)  Manganic  sulphate  Mn,(S04)3,  is 
soluble  in  dilute  sulphuric,  acid,  but  is  ivdiuvd  to  MnS04  by  the  attempt 
to  dissolve  it  in  water  alone;  potassium  manganic  sulphate  and  other 
manganic  alums  are  also  decomposed  by  water.  Alkali  manganates  and 
permanganates  are  soluble  in  water,  the  former  rapidly  changing  to  man- 
ganese dioxide  and  permanganate,  which  is  much  more  stable  in  solution. 
In  presence  of  reducing  agents  both  manganates  and  permanganates  are 
reduced  to  lower  forms. 

K2MnO4  +  8HC1  =  MnCl2  +  2KC1  +  2C12  +  4H2O 
2KMnO4  +  3MnSO4  -f  2H2O  =  5MnO2  +  K2SO4  +  2H,SO4 

Concentrated  H2S04  in  the  cold  dissolves  KMn04 ,  forming  (Mn03)2S0.1 
(a  sulphate  of  the  heptad  manganese:  2KMn04  +  3H2S04  =  (Mn03)2S04  + 
2KHS04  +  2H20  (Franke,  J.  pr.,  1887,  36,  31).  If  heat  be  applied  oxygen 
is  evolved  and  the  manganese  is  reduced  to  the  dyad  (4/). 

6.  Reactions,  a. — The  fixed  alkali  hydroxides  precipitate  from  solu- 
tions of  manganous  salts,  manganous  hydroxide,  Mn(OH)2 ,  white,  soon 
turning  brown  in  the  air  by  oxidation  to  manganic  hydroxide,  MnO(OH)  . 
The  precipitate  is  formed  by  the  reaction  of  the  negative  hydroxyl  ion  of 
the  alkali  with  the  positive  manganous  ion : 

2NH4  OH  +  Mn^f2  =  Mn(OH)2  +  2NH4  Cl . 

The  precipitate  is  insoluble  in  excess  of  the  alkalis  because  the  mangan- 
ous or  manganic  manganese  does  not  form  an  acid  ion.  Before  the 
manganous  manganese  is  oxidized  it  is  soluble  in  excess  of  ammonium 
salts  because  a  complex  salt  is  formed  in  which  the  manganese  forms  a 
part  of  the  acid  ion : 

+       -          ++ 4-      - 

Mn(OH)2  +  4NH4  Cl  -  (NH4)2  MnCU  +  2NH4  Cl . 

(1).  Ammonium  hydroxide  precipitates  one-half  of  the  manganese  as 
the  hydroxide  from  solutions  of  manganous  salts,  the  other  half  being 
held  in  solution  as  an  acid  ion  by  the  ammonium  salt  formed  (2)  (Da tu- 
rner, 3,  237).  The  presence  of  excess  of  ammonium  salt  prevents  the 
precipitation  of  the  manganese  by  ammonium  hydroxide  because  the 
manganese  is  present  in  the  acid  ion  (3)  (separation  of  manganese  from 
the  metals  of  the  third  group)  (Pickering,  J.  C.}  1879,  35,  672;  Lang- 
bein,  Z.,  1887,  26,  731).  Manganic  hydroxide,  MnO(OH),  is  insoluble 
in  the  alkalis  or  in  ammonium  salts.  It  gradually  precipitates,  com- 
pletely on  exposure  to  the  air,  as  a  dark  brown  precipitate  from  solutions 


180  MANGAXERE.  §134,  65. 

of  manganous  hydroxide  in  ammonium  salts.  Alkali  carbonates  pre- 
cipitate manganous  carbonate,  MnC03 ,  white,  oxidized  in  the  air  to  the 
brown  manganic  hydroxide,  and  before  oxidation,  somewhat  soluble  in  am- 
monium chloride.  Strong  ammonium  hydroxide  gradually  reduces  a  solu- 
tion of  potassium  permanganate  to  manganese  dioxide  (10ft). 

(1)  Mn  OH)2  +  4NH4C1  =  (NH4)2MnCl4  +  2NH4OH 
(*)  2MnS04  +  2NH40H  =  (NH4)2Mn(SO4)2  +  Mn(OH)2 
(3)  MnClz  +  2NH.C1  =  (NH4)2MnCl4 

&. — Oxalic  acid  and  alkaline  oxalates  precipitate  manganous  oxalate, 
soluble  in  mineral  acids  not  too  dilute.  All  compounds  of  manganese  of 
a  higher  degree  of  oxidation  are  reduced  to  the  manganous  condition  on 
warming  with  oxalic  acid,  or  oxalates  in  presence  of  some  mineral  acid: 
2KMn04  +  5H2C,04  +  3H2S04  =  K2S04  +  2MnS04  +  10C02  +  8H20  . 

Soluble  cyanides,  as  KCN  ,  precipitate  manganous  cyanide,  Mn(CN)2  ,  white, 
but  darkening-  in  the  air;  soluble  in  excess  of  the  precipitant  by  formation  of 
double  cyanides,  as  Mn(CN)2.2KCN  .  This  solution,  exposed  to  the  air,  pro- 
duces manga-nioyanides  (analogous  to  ferricyanides),  with  oxidation  of  the 
manganese:  12(Mn(CN)2.2KCN)  +  302  +  2H2O  =  8K8Mn(CN)a  +  4MnO(OH). 
Fe'"  and  Mn"  may  be  separated  by  treating  a  solution  of  the  two  metals  with 
a  strong  excess  of  KCN  and  then  with  iodine.  The  manganese  is  precipitated 
as  Mn02  and  the  iron  remains  in  solution  (Beilstein  and  Jawein.  B.,  187'J,  12, 
1528).  Ferrocyanides  piecipitate  white  inunyannux  fcrroci/anide,  Mn2Fe(CN)6  , 
soluble  in  hydrochloric  acid.  Ferricyanides  precipitate  brown  manganous  frrri- 
cyanide,  Mns(Fe(CN)6)2  ,  insoluble  in  acids  (separation,  with  Co  and  Ni  ,  from 
Zn)  (Tarugi,  Gazzetta,  1895,  25,  ii,  478).  If  an  alkali  or  alkali  carbonate  be 
present,  potassium  ferricyanide  oxidizes  manganous  compounds  to  manganese 
dioxide,  the  ferricyanide  being  reduced  to  ferrocyanide.  Potassium  ferro- 
cyanide  reduces  manganates  and  permanganates  to  manganous  compounds. 

c. — Nitric  acid  is  of  value  in  analysis  of  manganese  compounds  in  that 
it,  as  a  non-reducing  acid,  acts  readily  with  oxidizing  agents,  as  Pb02 , 
KC103 ,  etc.,  to  oxidize  manganous  compounds  to  manganese  dioxide  or  to- 
permanganic   acid.      Eeducing  agents   as   HC1 ,   etc.,   should  be   absent. 
Sulphuric  acid  may  be  used  instead  of  nitric  acid. 

2Mn(NO3)2  +  5Pb02  +  GHN03  =  2HMn04  +  5Pb(N03)2  +  2H2O 
5MnS04    f  2KC103  +  H2SO4  +  4H2O  =  5Mn02  +  K2S04  +  Cla  +  5H2SO4 

In  using  Pb02  and  HN03  to  detect  manganese,  the  compound  should  first 
be  reduced  w;th  hydrochloric  acid,  precipitated  with  potassium  hydroxide 
and  this  precipitate  dissolved  in  nitric  acid,  as  Mn02  is  not  all  oxidized 
by  Pb02  and  HN03  (Koninck,  Z.  angew.,  1889,  4). 

d. — ^Hypophosphorous  acid  reduces  all  higher  forms  of  manganese  to  the 
manganous  condition.  Alkali  phosphates,  as  Na,HPO4  ,  precipitate,  from 
neutral  solutions  of  manganous  salts,  normal  nia'ngamms  }>1iox}>1iate,  Mns(POjr)2  , 
white,  slightly  soluble  in  water,  and  soluble  in  dilute  acids.  It  turns  brown  in 
the  air.  The  manganous  hydrogen  phosphate,  MnHP04  ,  is  more  soluble  in 
water,  and  is  obtained  by  crystallization  from  a  mixture  of  manganous  sul- 


COLLCOfr 
Srf  PHARMACY 

§134, Qg.  MANGANESE.  181 

phate  acilulited  with  acetic  acid  and  di sodium  phosphate,  Na2HPO4 ,  added 
till  a  precipitate  begins  to  form.  From  the  ammonium-manganese  solution, 
freshly  formed  (6«),  phosphates  precipitate  nil  the  manganese  as  manyanoua 
ammonium  phosphate. 

e. — Hydrosulphuric  acid  precipitates  manganous  acetate  but  imperfectly, 
and  not  in  presence  of  acetic  acid,  and  does  not  precipitate  other  salts,  as 
manganous  sulphide  is  soluble  in  very  dilute  acids,  even  acetic  acid. 
Ammonium  sulphide  precipitates  from  neutral  solutions,  antl  forms  from 
the  recent  hydroxide  of  mixtures  made  alkaline,  the  flesh-colored  man- 
ganous sulphide,  MnS  .  Acetic  acid,  acting  on  the  precipitated  sulphides, 
separates  manganese  from  cobalt  and  nickel,  and  from  the  greater  part  of 
zinc.  All  the  higher  oxidized  forms  of  manganese  (in  solution  or  freshly 
precipitated)  are  reduced  to  the  manganous  condition,  with  separation  of 
sulphur  (10),  by  hydrosulphuric  acid  or  soluble  sulphides:  4KMn04  + 
14(NH4)JS  +  16H20  =:  4MnS  +  4KOH  +  28NH4OH  +  5S2 .  The  green 
manganous  sulphide,  MnS ,  crystalline,  anhydrous,  is  formed  by  the  action 
of  HoS  on  a  hot  ammoniacal  manganous  solution  not  containing  an  excels 
of  ammonium  salts  (Meineke,  Z.  angew.,  1888,  3),  also  by  pouring  the 
neutral  manganous  solution  into  a  hot  solution  of  ammonium  chloride 
and  excess  of  colorless  ammonium  sulphide.  The  fixed  alkali  sulphides 
produce  a  red  manganous  sulphide. 

Soluble  Sulphites  precipitate  from  solutions  of  manganous  salts,  manganous 
sulphite,  MnSO3  ,  white,  insoluble  in  water,  soluble  in  acids  (Gorgeu,  C.  r., 
1883,  96,  341).  Solutions  of  manganates  or  permanganates  are  immediately 
reduced  to  the  flocculent  brown-black  manganese  dioxide  by  solutions  of  sodium 
sulphite  or  sodium  thiosulphate ;  if  acids  be  present,  the  reduction  is  complete 
to  manganous  salts. 

/. — HC1 ,  HBr ,  and  HI  readily  reduce  the  higher  compounds  of  man- 
ganese to  manganous  salts  with  evolution  of  the  corresponding  halogen. 
When  manganese  dioxide  is  dissolved  in  concentrated  HC1  without  heat, 
the  dark  brownish  colored  solution  contains  manganese  tetraohloride, 
MnCl4 ,  and  trichloride,  MnCl3  which  deposits  Mn02  on  dilution  with 
water  and  on  warming  decomposes  into  manganous  chloride  and  chlorine 
(56)  (Pickering,  J.  C.,  1879,  35,  654,  W.  B.  Holmes,  J.  Am.  Soc.  29, 
1277).  Potassium  iodide  instantly  reduces  a  solution  of  potassium  per- 
manganate, forming  manganese  dioxide  and  an  iodate  (distinction  from 
chloride  and  bromide).  Potassium  chlorate  or  bromate  when  boiled  with 
concentrated  nitric  or  sulphuric  acids  and  manganous  compounds  forms 
manganese  dioxide  (c). 

g. — Soluble  arsenites  precipitate  manf/anous  arsenite,  and  arsenates  precipitate 
HHUIWIIHHIH  <(rw>Hit<\  insoluble  in  water,  soluble  in  acids.  Arsenous  acid  and 
arsenites  reduce  solutions  of  manipulates  or  permanganates,  forming-  a  brown 
flocculent  precipitate1:  or  a  colorless  solution  if  warmed  in  presence  of  a 
mineral  acid.  //.Normal  potassium  chromate  precipitates  manganous  salts, 
brown,  soluble  in  acids  and  in  ammonium  hydroxide;  no  precipitate  is  formed 
with  potassium  dichroinate.  «.— Soluble  manganates  and  permanganates  pre- 


182  MANGANESE.  §134,  7. 

cipitate  manganous  salts  as  manganese  dioxide,  being  themselves  reduced  to  the 
h  ame  form:   3MnSO4  +  2KMnO4  +  2H2O  =  5MnO2  +  K2SO4  +  2HSSO4  . 

7.  Ignition  with  alkali  and  oxidizing  agents,  forming  a  bright  green  mass 
of  alkaline  manganate,  constitutes  a  delicate  and  convenient  test  for  man- 
ganese, in  any  combination.  A  small  portion  of  precipitate  or  fine  powder 
is  taken.  If  the  manganese  forms  but  a  small  part  of  a  mixture  to  be 
tested,  it  is  better  to  submit  the  substance  to  the  systematic  course  of 
analysis,  and  apply  this  test  to  the  precipitate  by  alkali,  in  the  fourth 
group.  A  convenient  form  of  the  test  is  by  ignition  on  platinum  foil  with 
potassium  or  sodium  nitrate  and  sodium  carbonate  (a).  Ignition,  by  an 
oxidizing  flame,  on  platinum  foil,  with  potassium  hydroxide,  effects  the 
same  result,  less  quickly  and  perfectly  (&).  Ignition  by  the  oxidizing  flame 
of  the  blow-pipe,  in  a  bead  of  sodium  carbonate,  on  the  loop  of  platinum 
wire,  also  gives  the  green  color  (c). 
(a)  Mn(OH)2  +  2KNO3  +  NjfcCO  3=  .NasMnO,  +  2KNO2  +  CO2  +  H2O 

or  if  a  small  amount  of  KNOs  is  present, 
(fe)     3Mn(OH)3  +  4KNO3  +  NaaCOs  =2K2MnO4  +  Na^MnCX  +  4NO  +  CO2  +  3H2O 

o       Mn  OH)2  +  2KOH  +  O2  =  K2MnO4  +  2H2O 

,1      MnvOH)2  +  NaaCOs  +  O2  =  NaaMnO,  +  H2O  +  CO2 


With  beads  of  borax  and  microcosmic  salt,  before  the  outer  blow-pipe  flame, 
manganese  colors  the  bead  violet  while  hot,  and  amethyst-red  when  cold.  The 
color  is  due  to  the  formation  of  manganic  oxide,  the  coloring-  material  of  the 
amethyst  and  other  minerals,  and  is  slowly  destroyed  by  application  of  the 
inner  flame,  which  reduces  the  manganic  to  manganous  oxide. 

8.  Detection.  —  After  the  removal  of  the  metals  of  the  first  three  groups 
(the  third  group  in  the  presence  of  NH4C1  in  excess,  5fr  and  6a),  the  Mn 
with  Co  ,  Ni  and  Zn  is  precipitated  in  the  ammoniacal  solution  by  H2S  . 
By  digestion  in  cold  dilute  HC1  the  sulphides  of  Mn  and  Zn  are  dissolved, 
and  after  boiling  to  remove  the  H2S  ,  Mn  is  precipitated  as  the  hydroxide 
by  excess  of  KOH  ,  which  dissolves  the  Zn  .     The  precipitate  of  the  man- 
ganese is  dissolved  in  HN03  and  boiled  with  more  HN03  and  an  excess  of 
Pb02  .     A  violet-colored  solution  is  evidence  of  the  presence  of  manganese. 

9.  Estimation.  —  (1)   By  converting  into  Mn304    (4c),   and  weighing  as  such. 

(2)  By  precipitating  as  MnNH4P04  ,  and  after  ignition  weighing  as  Mn2P2O7  . 

(3)  By  treating  the   neutral   manganous   salt   with   a   solution    of   KMnO4    of 
known  strength  (Qi).     If  some  ZnS04  is  added  the  action  is  more  satisfactory 
(Wright  and  Menke,  J.  C.,  1880,  37,  42).     (Jf)  By  boiling  the  manganous  com- 
pound with  Pb02  and  HNO3  ,  and  comparing  the  color  with  a  permanganate 
solution  of  known  strength   (Peters,  C.  N.,  1870,  33,  :J5).     (5)   The  manganous 
compound    is   oxidixed    to   MnO2    by   boiling  with   KC103    and   HNO3  .     This  is 
then  reduced  by  an  excess  of  standard  H2O,  ,  H,,C,O4  or  FeSO4  ,  and  the  excess 
of  the  reagent  estimated  by  the  usual  methods.     (6)  MnO,  ,  obtained  as  in  (.5), 
is  treated  with  H2C204  and  the  evolved  CO2  measured  or  weighed.     (7)  Mn02  , 
obtained  as  in  (o),  is  boiled  with  HC1  and  the  evolved  Cl  estimated. 

10.  Oxidation.—  (a)  Mn"   is  oxidized  to  Mn'"  in  alkaline  mixture  on 
exposure  to  the  air  ;  to  Mnlv  in  neutral  solution  by  KaMn04  and  KMn04  , 


§135,  3.  ZINC.  183 

in  alkaline  mixture  by  Cl ,  Br ,  I ,  K,Fe(CN)0 ,  KC10 ,  KBrO ,  H.O,1,  etc.; 
in  acid  solution  by  boiling  with  concentrated  HN03  or  H2S04 ,  arid  KC103 
or  KBrO, .  MnVI-n  is  oxidized  to  MnVI  by  fusion  with  an  alkali  and  an 
oxidizing  agent,  or  by  fusion  with  KC103  alone  (Boettger,  Z.,  1872,  11, 
433).  MnVII~n  is  oxidized  to  Mnvn  by  warming  with  Pb02  or  Pb304  and 
HN03  or  H2S04 .  The  higher  oxide  of  lead  should  be  in  excess  and  reduc- 
ing agents  should  be  absent  as  they  delay  the  reaction;  hence  in  analysis 
the  manganese  should  be  precipitated  as  the  hydroxide  or  sulphide,  fil- 
tered, washed,  and  then  dissolved  in  HN03  or  H2S04 ,  and  boiled  with  the 
higher  oxide  of  lead  (6c).  A  solution  of  potassium  manganate  decomposes 
into  potassium  permanganate  and  manganese  dioxide  on  standing,  more 
rapidly  on  warming  or  dilution  with  water,  (b)  All  compounds  of  man- 
ganese having  a  higher  degree  of  oxidation  than  the  dyad,  (Mn"+n)  are 
reduced  to  the  dyad  (Mn")  by  H2C204 ,  HH2P02 ,  H2S4,  K2S ,  H2SO, ,  H2022 
(in  neutral  or  alkaline  solution  to  MnIV),  HC1 ,  HBr ,  HI ,  HCNS  ,  Hg',  Sn", 
As'",  Sb'",  Cu',  Fe",  Cr",  Cr'",  etc.;  the  reducing  agents  becoming  respec- 
tively C02 ,  Pv,  S°  to  SVI  (depending  upon  the  temperature,  concentration, 
and  the  agent  used  in  excess),  Cl ,  Br ,  I ,  HCN  and  SVI,  Hg",  SnIV,  Asv,  Sbv, 
Cu",  Fe'",  and  CrVI.  MnIV+n  is  reduced  to  MnIV  (or  Mn'")  by  H  3,  Asfl.,3, 
SbH33,  PH,3,  Na2S034,  Na2S2034,  NH4OH3  (slowly),  Mn",  etc.  KMn04  is 
reduced  to  K2Mn04  on  boiling  with  concentrated  KOH  :  4KMn04  +  4KOH 
=  4K,Mn04  +  2H20  +  02  (Eammelsberg,  £.,  1875,  8,  232). 
§135.  Zinc.  Zn  =  65.37.  Valence  two. 

1.  Properties.—  Specific  gravity,   7.142    (Spring,   B.,   1883,  16,  2723).     Melting 
point,   418.5°   to   419.35°    (Burgess,    Wash.   Acad.   of  Sci.,    1-18).     Boiling  point, 
918°  (Berthelot,  C.  r.,  1912,  134).     It  is  a  bluish-white  metal,  retaining  its  lustre 
in  dry  air,  but  slightly  tarnished  in  moist  air  or  in  water.     When  heated  to  the 
boiling  point  with  abundant  excess  of  air  it  burns  with  a  bluish-white  flame  to 
zinc   oxide.      Zinc   dust   mixed  with  sulphur  is  ignited  by  percussion    (Schwarz, 
B.,  1882,  16,  2505).     At  ordinary  temperature  it  breaks  with  a  coarse  crystalline 
fracture.     It  is  more  malleable  at  100°  to  150°  than  at  other  temperatures,  and 
at  that  temperature  may  be  drawn  into  wire  or  rolled  into  sheets.     At  205°  it 
is  so  brittle  that  it  may  be  easily  powdered  in  a  mortar. 

Zinc  finds  an  extended  use  in  laboratories  for  the  generation  of  hydrogen. 
It  is  molded  in  sticks  or  granulated  by  pouring  the  molten  metal  into  cold 
water.  The  pure  metal  is  not  suitable  for  the  generation  of  hydrogen,  as  the 
reaction  with  acids  proceeds  too  slowly  (Weeren,  B.,  1891,  24,  1785).  Com- 
mercial impurities  render  the  metal  readily  soluble  in  acids,  or  the  pure  metal 
may  be  treated  with  a  dilute  solution  of  platinum  chloride  (twenty  milligrams 
PtCl4  per  litre),  or  copper  sulphate.  Met.-illic  platinum  or  copper  is  deposited 
upon  the  zinc:  PtCl,  +  2Zn  =  Pt  +  2ZnCL  ,  CuSO4  +  Zn  =  Cu  +  ZnSO4. 

2.  Occurrence. — It  is  found  as  calamine  (3nCO3),  as  zinc-blende    (ZnS);    also 
associated  with  other  metals  in  numerous  <>  es. 

3.  Preparation. — The   process  usually  employed    consists   of  two   operations: 

'Klein,  Arch.  Pharm.,  1880,  227,  77;  Jannaesch  and  von  Cloedt,  Z.  annrg.,  1805,  10,  398  and  410; 
Carnot,  C.  r.,  1888,  1O7,  997  and  1150. 

*Carnot,  Bl.,  1889,  (3),  1,  277  ;  Gorgeu,  C.  r.,  1800,  HO,  958.  3  Jones,  J.  C.,  1878,  33,  96.  4  Hoenig 
and  Zateck,  Jf.,  1863,  4,  738  ;  Glaeser,  Jf.,  1685,  6,  329. 


184  ziNC.  §135,  4. 

(I)  Roasting-:  in  case  of  the  carbonate  the  action  is:  ZnCO,  =  ZnO  -f  CO,-  if  it 
is  a  sulphide,  2ZnS  +  3O2  =  2ZnO  +  2SO2  .  (2)  Reduction  with  distillation; 
after  mixing-  the  ZnO  with  one-half  its  weight  of  powdered  coal,  it  is  distilled 
at  a  white  heat.  Its  usual  impurities  are  As,  Cd  ,  Pb  ,  Cu  ,  Fe  and  Sn  .  It  is 
purified  by  repeated  distillation,  each  time  rejecting-  the  first  portion,  which 
contains  the  more  volatile  As  and  Cd  ,  and  the  last  which  contains  the  less 
volatile  Pb  ,  Cu  ,  Fe  and  Sn  .  Strictly  chemically  pure  zinc  is  best  prepared 
from  the  carbonate  which  has  been  purified  by  precipitation. 

4.  Oxide  and  Hydroxide. — Zinc  oxide   (ZnO)   is  made  by  igniting-  in  the  air 
either  metallic  zinc,  its  hydroxide,  carbonate,  nitrate,  oxalate,   or  any  of  its 
org-anic  oxysalts.     Zinc  hydroxide,   Zn(OH),  ,  is  made  from  solutions   of   zinc 
salts  by  precipitation  with  fixed  alkalis  (6a). 

5.  Solubilities.— (a)  Metal.— Pure  zinc  dissolves  very  slowly  in  acids  or  alkalis, 
unless  in  contact  with  copper,  platinum  or  some  less  positive  metal    (Baker, 
J.   C.,   1885,  47,   349).     The   metallic   impurities   in   ordinary   zinc   enable   it    to 
dissolve   easily  with   acids   or  alkali    hydroxides.     In   contact   with   iron,   it   is 
quite   rapidly   oxidized   in   water   containing1  air,   but   not   dissolved   by    water 
unless  by  aid  of  certain  salts.     It  dissolves  in  dilute  hydrochloric,   sulphuric  * 
and  acetic  acids  (Jf),  and  in  the  aqueous  alkalis  (2),  with  evolution  of  hydrogen; 
in  very  dilute  nitric  acid,  without  evolution  of  gas   (3);  in  moderately  dilute 
cold  nitric  acid,  mostly  with  evolution  of  nitrous  oxide   (j);  and,  in  somewhat 
less  dilute  nitric  acid,  chiefly  with  evolution  of  nitric  oxide  (5).     Concentrated 
nitric  acid  dissolves  zinc  but  slightly,  the  nitrate  being-  very  sparingly  soluble 
in  nitric  acid   (Montemartini,   Gazzstta,  1892,  22,  277).     Hot  concentrated  sul- 
phuric acid  dissolves  it  with  evolution  of  sulphur  dioxide  (6). 

(1)  Zn  +  H2S04  =  ZnS04  +  H2 

(2)  Zn  +  2KOH  =  K2Zn02  +  H2 

(3)  4Zn  +  10HN03  =  4Zn(N03)2  +  NH4NO3  +  3H2O 

(4)  4Zn  +  10HN03  =  4Zn(N08),  +  N20  +  5H20 

(5)  3Zn  +  8HN03  =  3Zn(N08)2  +  2NO  -f  4H20 

(6)  Zn  +  2H2S04  =  ZnS04  +  S02  +  2H20 

(&)  Oxide  and  Hydroxide. — All  the  agents  which  dissolve  the  metal,  dissolve  also 
its  oxjde  and  hydroxide. 

(c)  Salts. — The  chloride,  bromide,  iodide,  chlorate,  nitrate  (6aq),  and 
acetate  (7aq)  are  deliquescent;  the  sulphate  (7aq)  is  efflorescent.  The 
chloride  is  readily  soluble  in  alcohol  in  all  proportions  (Kremers,  Pogg., 
1862,  115,  360).  The  sulphide,  basic  carbonate,  phosphate,  arsenate, 
oxalate,  and  ferrocj^anide  are  insoluble  in  water;  the  sulphite  is  sparingly 
soluble.  The  ferrocyanide  is  insoluble  in  hydrochloric  acid  (Fahlberg,  Z., 
1874,  13,  380).  The  sulphide  is  almost  insoluble  in  dilute  acetic  acid  (sepa- 
ration from  MnS).  All  zinc  salts  are  soluble  in  KOH  and  NaOH  except 
zinc  sulphide,  and  all  in  NH4OH  except  ZnS  and  Zn2Fe(CN)6 . 

6.  Reactions,    a. — The  fixed  alkali  hydroxides  precipitate  zinc  hydroxide, 
Zn^OH)2 ,  white,  soluble  in  excess  of  the  precipitant  forming  an  alkali  ziticate: 

ZnCl2  +  2KOH  =  Zn(OH)2  +  2KC1 
Zn(OH)2  +  2  KOH  =  K2ZnO2  +  2H2O 

Ammonia  precipitates  from  neutral  solutions  free  from  ammonium  salts, 
zinc  hydroxide,  soluble  in  excess  of  ammonia  or  ammonium  salts  forming 
complex  zinc  ammonia  ions: 

ZnCl2  +  2NH4OH  =  Zn(OH)2  +  2NH4C1 
Zn(OH)2  +  6NH3  =  Zn(NH3)6(OH)2 
*  Muir  and  Jlobbs,  C.  A'.,  1882.  45,  69. 


§135,  6A.  ZINC.  185 

The  precipitate  of  zinc  hydroxide  dissolves  more  readily  in  excess  of  the 
alkalis  at  ordinary  temperature  than  when  heated.  Unless  a  strong  excess 
of  the  alkali  be  present,,  boiling  causes  a  precipitation  of  zinc  oxide,  more 
readily  from  the  solution  in  ammonium  hydroxide  than  in  the  fixed 
alkalis.  The  presence  of  other  metals — as  iron  or  manganese — makes 
necessary  the  use  of  much  more  alkali  to  effect  solution.  An  alkali  solu- 
tion as  dilute  as  tenth  Normal' does  not  dissolve  zinc  hydroxide,  no  matter 
how  great  an  excess  be  added  (Prescott,  J.  Am.  Soc.,  1880,  2,  29). 

Alkali  carbonates  precipitate  the  basic  carbonate,  Zn5(OH)6(C03)2 ,  white, 
soluble  in  ammonium  carbonate,  readily  in  alkali  hydroxides  (Kraut,  Z. 
anorg.,  1896,  13,  1).  Carbonates  of  Ba,  Sr,  Ca ,  and  Mg  have  no  action 
at  ordinary  temperatures  (separation  from  Fe'",  Al ,  and  Cr'"),  but  upon 
boiling  precipitate  the  whole  of  the  zinc. 

ft. — Alkali  cyanides,  as  KCN ,  precipitate  zinc  cyanide,  Zn(CN)2  ,  white, 
soluble  in  excess  of  the  precipitant.  Alkali  ferrocyanides,  as  K4Fe(CN)R  , 
precipitate  zinc  ferrocyanide,  Zn2Ee(CN)a  ,  white  (5c).  Alkali  ferricyanides, 
as  K3Pe(CN)6,  precipitate  zinc  ferricyanide,  Zn3(Fe(CN)0)2 ,  yellowish,  c. — 
See  5c.  ^.—Sodium  phosphate,  Na2HPO4  ,  precipitates  zinc  phosphate,  soluble 
in  alkali  hydroxides  and  in  nearly  all  acids. 

e. — Hydrosulphuric  acid  precipitates  a  part  of  the  zinc  from  neutral 
solutions  of  its  salts  with  mineral  acids,  and  the  whole  from  the  acetate; 
also  from  other  salts  of  zinc,  by  addition  of  alkali  acetates  or  monochlor- 
acetic  acid,  in  small  excess  (separation  from  Mn ,  Co ,  Ni ,  and  Fe)  (Berg, 
Z.,  1886,  25,  512):  ZnCl2  +  2KC2H302  +  H2S  =  ZnS  +  2KC1  + 
2HC2H302  .*  That  is:  Zinc  sulphide  is  not  entirely  soluble  in  dilute  acids, 
though  much  more  soluble  in  mineral  acids  than  in  acetic  acid.  The 
precipitate  is  white  when  pure.  Alkali  sulphides  completely  precipitate 
zinc  as  sulphide,  both  from  its  salts  with  acids  and  from  its  soluble  com- 
binations with  alkalis. 

Concentrated  solutions  of  sodium  sulphite  precipitate  solutions  of  zinc  salts 
as  basic  zinc  sulphite;  or  if  the  solutions  be  too  dilute  for  immediate  precipita- 
tion, boiling  will  cause  the  immediate  formation  of  the  bulky  white  precipitate 
of  the  basic  sulphite  (Seubert,  Arch.  Pharm.,  1891,  229,  316).  f.—If  a  hot  con- 
centrated zinc  chloride  solution  be  treated  with  ammonium  hydroxide  until 
a  precipitate  begins  to  form,  a  basic  chloride,  2ZnCl2.9ZnO  ,  will  separate  out 
upon  cooling  as  a  white  precipitate  (Habermann,  M.,  1884,  5,  432). 

g. — Zinc  salts  are  precipitated  by  solutions  of  alkali  arsenites  and  arsenates, 
forming  respectively  zinc  arsenite  or  arsenate,  white,  gelatinous,  readily  solu- 
ble in  alkalis  and  acids,  including  arsenic  acids,  ft.— Normal  potassium  chro- 

*  In  the  equation  for  acetic  acid,  ab  =kc,  a  and  b,  the  concentrations  of  the  H  and  C,II3O, 

ions  respectively,  are  small,  c  is  large,  and  k,  the  so-called  "dissociation-constant,"  to  which 
the  strength  of  the  acid  is  proportional,  is  very  small.  But  addition  of  the  fully-dissociated 
sodium  acetate  to  the  likewise  completely-ionized  hydrochloric  acid  gives  a  solution  containing 
the  ions  in  very  large  concentration  and  practically  none  of  the  non-dissociated  acetic  acid. 
To  restore  equilibrium  the  H  ions  of  the  HC1  unite  with  the  acetic  ions  of  the  sodium  acetate, 
leaving  Na  and  Cl  ions  in  the  solution.  The  displacement  of  a  weak  acid  from  its  salt  by  a 
strong  one  lies  then  not  so  much  in  an  attraction  of  the  strong  acid  by  the  base  as  in  the  ten- 
dency of  the  weak  acid  to  form  the  non-ionized  molecule. 


186  ZINC.  §135,  7. 

mate  forms,  with  solutions  of  zinc  salts,  a  yellow  precipitate  readily  soluble 
in  alkalis  and  acids,  including  chromic  acid.  No  precipitate  is  formed  with 
K2Cr207  . 

7.  Ig-nition. — With  sodium  carbonate,  on  charcoal,  before  the  blow-pipe,  com- 
pounds of  zinc  are  reduced  to  the  metallic  state.     The  metal  is  vaporized,  and 
then  oxidized  in  the  air,  and  deposited  as  a  non-volatile  coating-,  yellow  when 
hot  and  white  when  cold.     If  this  coating-,  or  zinc  oxide  otherwise  prepared, 
be  moistened  with  solution  of  cobalt  nitrate  and  again  ignited,  it  assumes  a 
green  color  (Bloxam,  J.  C.,  1865,  18,  98).     With  borax  or  microcosmic  salt,  zinc 
compounds  give  a  bead  which,  if  strongly  saturated,  is  yellowish  when  hot, 
and  opaque  white  when  cold. 

8.  Detection. — After  the  removal  of  the  first  three  groups,  the  Zn  is 
precipitated  with  Co ,  Ni  and  Mn  from  the  ammoniacal  solutions  by  H2S . 
Digestion  of  the  precipitated  sulphides  with  cold  dilute  HC1  dissolves  the 
Mn  and  Zn  as  chlorides.     The  solution  is  thoroughly  boiled  to  expel  the 
H2S  and  the  zinc  changed  to  Na2Zn02  by  an  excess  of  NaOH  ,  which  precipi- 
tates the  manganese  as  the  hydroxide.  From  the  alkaline  filtrate  H2S  gives 
a  white  or  grayish-white  precipitate — evidence  of  the  presence  of  Zn  . 

9.  Estimation. — (1)    Zinc    is    weighed    as    an   oxide,    into   which    form    it   is 
brought   by   simple    ignition   if   combined   with   a  volatile    inorganic    oxyacid, 
otherwise  it  should  be  changed  to  a  carbonate  and  then  ignited.     (2)   It  is 
converted  into  a  sulphide,  and  after  adding  powdered  sulphur  it  is  ignited  in 
a  stream  of  hydrogen  or  hydrogen  sulphide,  and  weighed  as  a  sulphide  (Kiinzel, 
Z.,  1863,  2,  373).     (3)   It  may  be  converted  into  ZnNH4P04 ,  and,  after  drying 
at  100°,  weighed.     Ignition  converts  it  into  Zn,P,07  ,  with  slight  loss  of  zinc. 
(4)   Volumetrically,  by  converting  into  Zn_,Fe(CN)6  and  titrating  with  potas- 
sium permanganate  or  by  using  FeCl3   acidulated  with  HC2H302  as  external 
indicator  (Voigt,  Z.  angew,  1889,  307).     (5)  By  precipitation  as  Zn3(Fe(CN)ft,)2  , 
treating  the  precipitate  with  potassium  iodide  and  titrating  the  liberated  iodine 
(Mohr,  DingL,   1858,  48,   115).     (6')   By   titration   in  hydrochloric  acid   solution 
with  K4Fe(CN)6  ,  using  a  uranium  salt  as  an  indicator  (Fahlberg,  Z.,  1874,  13, 
379;   Koninck   and   Prost,   Z.    angew.,   1896,   568).     (?)    By   titration  in    alkaline 
solution  with  Na2S  ,  using  a  copper  salt  as  an  indicator.     (8)  The  zinc  is  pre- 
cipitated as  ZnNH4AsO4  ,  the  precipitate  decomposed  with  HI  and  the  liber- 
ated iodine  titrated  with  standard  Na2S2O3  (Meade,  J.  Am.  Soc.,  1900,  22,  353). 

10.  Oxidation. — Metallic  zinc  precipitates  the  free  metal  from  solutions 
of  Cd ,  Sn ,  Pb ,  Cu ,  Bi ,  Hg ,  Ag ,  Pt ,  An ,  As ,  Sb ,  Te ,  In ,  Fe  *,  Co , 
Ni,  Pd,  Rh,  Ir,  and  Os  (Gmelin-Kraut,  Handbuch,  1875,  3,  6).     Zinc 
with  copper  (zinc-copper  couple,  used  in  water  analysis)  reduces  nitrates 
and  nitrites  to  ammonia,  chlorates  to  chlorides,  iodates  to  iodides,  ferri- 
cyanides  to  ferrocyanides,  etc.  (Thorpe,  /.  (7.,  1873,  26,  541).     Solutions 
of  chromates  are  reduced  to  chromic  salts,  ferric  salts  to  ferrous  salts, 
and  compounds  of  manganese  having  more  than  two  bonds  are  reduced  to 
the  dyad  in  presence  of  some  non-reducing  acid.     Zinc  is  precipitated  as 
the  metal  from  acetic  solutions  by  Mg  (Warren,  C.  N.,  1895,  71,  92). 
The  oxide  is  reduced  to  the  metal  by  heating  in  a  current  of  hydrogen 
(Deville,  A.  Ch.,  1855  (3),  43,  477). 

*Davie8.  J.  C.,  1875,  28, 3a. 


§136. 


REACTIONS  OF  IRON  AND  ZINC  GROUP  BASES. 


187 


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TABLE  FOR  ANALYSIS  OF   THE  ZINC  G&OUP. 


§131 


§137.  TABLE  FOE  ANALYSIS  OF  THE  ZINC  GROUP  (FOURTH  GROUP) 
(Phosphates  and  Oxalates  being  absent). 

Into  the  clear  ammoniacal  filtrate  from  the  Third  Group  pass  Hydrogen  Sul- 
phide, and  if  a  precipitate  appears,  warm  until  it  subsides.  Filter  and  wash 
with  a  one  per  cent  solution  of  NH4C1  .  (Test  nitrate,  in  which  H2S  gives 
no  precipitate  for  the  Fifth  Group.) 

Precipitate:  CoS  ,  NiS  ,  MnS  ,  ZnS  . 

Treat  on  the  filter  with  cold  dilute  Hydrochloric  Acid  (1-4). 


Residue:  CoS,  NiS*  (black). 

Test  with  the  borax  bead.  A  blue 
bead  indicates  cobalt,  (§132,  7). 

Dissolve  the  remainder  of  the  sulphides 
in  nitro-hydrochloric  acid  or  HC1  + 
crystal  of  KC1O3  .  Evaporate  to  ex- 

S^l   excess   of    CL_  ,    neutralize   with 
H4OH  and  divide  into  two  parts. 


Solution:  MnCl2  ,  ZnCl2(H2S,HCl). 
(traces  of  CoCL>  and  NiCl2.) 

Boil  the  solution  thoroughly  to  remove  the 
H2S  ,  cool,  and  add  a  decided  excess 
of  potassium  or  sodium  hydroxide  and 
bromine  water  and  heat  (§135,  6a). 
Filter  and  wash. 


For  Cobalt: 

For  Nickel: 

Precipitate  : 

Solution  : 

Add   NaHCO3   and 
H2O2  ;      warm 

Add  dimethyl-gly- 
oxime  and  warm. 

MnO2  (traces  of  Co(OH)3  and 

Ni(OH)3  . 

K2ZnO2 

gently  and  filter. 

A  scarlet  precipi- 

Dissolve in  HNO3  -f  a  small 

Test     for 

A  green  color  to 

tate  shows  nickel. 

amount  of  H2O2  and  boil. 

zinc    by 

the  filtrate  indi- 

(If Co  and  Ni  were  present 

adding 

cates     cobalt 

Or:   Add  excess  of 

precipitate  the  manganese 

H2S  .       A 

(§140). 

hot  KOH  and  Br, 

from     this     solution     by 

white  pre- 

boil,  filter,  wash 

NH4OH+H2O2  and  boiling. 

cipitate 

If  sufficient  nickel 

(until     fi  1  1  r  a  t  e 

Filter,  wash  and  redissolve 

(ZnS)    in- 

be present  to  ob- 

gives no  precipi- 

in HNO3+H2O2  and  boil.) 

dicates 

scure     the     blue 

tate  with  AgNOs), 

Boil  with  HNO3  and  Pb.C^ 

zinc. 

bead,  add  to  the 

add    solution    of 

or  PbO2.     Manganese  will 

solution    of    the 

hot  KI  and   test 

give  the  reddish  purple  color 

sulphides    (§133, 

the  filtrate  with 

ofHMnO4(§134,6c).  Man- 

7) an  excess  of  ni- 

CS2.     If  free  io- 

ganese may  also  be  tested 

troso-/3-naphthol 

dine    appears, 

for,by  pouring  diamine  over 

in  acetic  acid  so- 

nickel is  present 

the  Mn02  if  Co  and  Ni  are 

lution  (§132,  66); 

(§133,  6/). 

absent,  or  if  they  are  pres- 

filter, wash,  and 

ent,  by  pouring  it  over  the 

test  the  brick-red 

precipitate    of    Mn(OH)3 

precipitate    with 

obtained  by  precipitating 

the  borax  bead. 

with     NH4OH  +  H2O2  . 

Manganese  gives  a  reddish 

The  filtrate  may  be 

purple  precipitate. 

tested  for  nickel. 

Dark-colored  original  solu- 

tions indicating  an  alkali 

salt  of  manganese  should 

be    reduced    by    warming 

Study     the 

Study      §132,      6c, 
§135,  §138,  §139, 

Study  the  text  at 
§133,  6a,  6,  e  and 

with  HC1  before  proceed- 
ing with  the  analysis  (§134, 
5c  and  6/). 

text   at 
§135,     6a 

and   e, 

§140,  §141,  §144, 

/;  §132   66  and  c~ 

§136,  §138, 

§135  and  ff. 

§136,  §138,  §139, 

Confirm  by  study  of  the  text 

§139,  §142, 

§140,  §141,  §144, 

§134,  7,  §136,  §138,  §139, 

§143,  §144, 

§145  and  ff. 

§142,  §143,  §144,  §145  and 

§145  and/. 

ff- 

*  Small  portions  of  cobalt  and  nickel  sulphides  may  be  dissolved  by  the  cold  dilute  HC1, 
and  will  be  precipitated  with  the  Mn(OH)2 . 


§140.  DIRECTIONS  FOR  ANALYSIS   WITH  NOTES.  189 

DIRECTIONS  FOR  THE  ANALYSIS  OF  THE  METALS  OF  THE  FOURTH  GROUP. 
§138.  Manipulation.— Into  the  warm  strongly  a  mm  onincal  filtrate  from  the 
third  group  (§128) ,  H^S  gas  is  passed  until  complete  precipitation  is  obtained: 
MnCL.2NH4Cl  +  2NH4OH  +  H2S  =  MnS  +  4NH4C1  +  2H2O 
(NH4)2ZnO2  +  2H2S  =  ZnS  +  (NH4)2S  +  2H2O 

The  solution  is  warmed  until  the  precipitate  subsides,  allowed  to  stand 
for  a  few  minutes,,  and  is  then  filtered  and  the  precipitate  washed  with 
hot  water  containing  about  one  per  cent  of  NH4C1  (§139,  2).  The  filtrate 
should  be  again  tested  with  H2S  and  if  complete  precipitation  has  been 
obtained  it  is  set  aside  to  be  tested  for  the  metals  of  the  succeeding  groups 
(§191).  The  well  washed  precipitate  of  the  sulphides  of  Co  ,  Ni ,  Mn ,  and 
Zn  is  digested  on  the  filter  or  in  a  test-tube  with  cold  dilute  H£l  (one  part 
of  reagent  HC1  to  four  of  water) :  MnS  +  2HC1  =  MnCl,  -f  H2S .  The 
black  precipitate  remaining  undissolved  contains  the  sulphides  of  Co  and 
Ni,  the  filtrate  contains  Mn  and  Zn  as  chlorides  with  an  excess  of  HC1 
and  the  H2S  which  has  not  escaped  as  the  gas. 

§139.  Notes. — (1)  Instead  of  passing1  the  H2S  into  the  ammoniacal  solution,  a 
freshly  prepared  solution  of  ammonium  sulphide  may  be  used.  The  yellow 
ammonium  sulphide,  (NH4)2SX  ,  should  not  be  employed  to  precipitate  .the 
metals  of  the  fourth  group,  as  nickel  sulphide  is  quite  appreciably  soluble  in 
that  reagent  (§133,  6e). 

(2)  The  sulphides  of  the  fourth  group,  especially  MnS  and  ZnS  ,  should  not 
be  washed  with  pure  water,  as  they  may  be  changed  to  the  colloidal  sulphides, 
soluble  in  water.     The  presence  of  a  small  amount  of  NH4C1  prevents  this,  and 
does  not  in  any  way  interfere  with  the  analysis  of  the  succeeding  groups. 

(3)  If  the  precipitates  are  to  be  treated  on  the  filter  with  the  dilute  HC1, 
the  acid  solution  should  be  poured  on  the  precipitate  three  or  four  times.     For 
digestion  in  a  test  tube,  the  point  of  the  filter  is  pierced  and  the  precipitate 
washed  into  the  test  tube  with  as  little  water  as  possible. 

(4)  The  sulphides  of  Co  and  Ni  are  not  entirely  insoluble  in  the  cold  dilute 
HC1 ,  and  traces  of  them  may  usually  be  detected  in  the  precipitate  for  Mn 
(§137,  footnote). 

(5)  Dilute  acetic  acid  readily  dissolves  MnS  but  scarcely  attacks  ZnS  (§135, 
6e).     If  desired,  dilute  acetic  may  be  used,  first  removing  the  Mn  and  then 
adding  dilute  HC1  to  dissolve  the  Zn  . 

(6)  If  large  amounts  of  iron  are  present,  a  portion  of  the  Mn  will  always 
appear  in  the  third  group  (§134,  6fl),  and  is  detected  by  the  green  color  of  the 
fused  mass  when  testing  for  Cr:  3Mn(OH)2  +  4KNO3  -f-  Na2CO3  =  2K,MnO4  + 
Na2MnO4   +  4NO  -f  CO,   -f  3TL..O  .     Too  much  HNO3  in  the  oxidation  of  the 
iron  favors  this  precipitation  of  Mn  with  Fe'"  due  to  the  oxidation  of  the  Mn  to 
the  triad  or  tetrad  combination. 

(7)  Small  amounts  of  the  Fifth  Group  elements  are  carried  down  with  the 
ammonium  sulphide  precipitate.     As  much  as  5  mg.  of  barium  may  be  present 
in  this  precipitate  (Curtman  &  Frankel,  J.  Am.  Soc.,  33,  724,  1911). 

§140.  Manipulation. — The  black  precipitate  of  cobalt  and  nickel  sul- 
phides should  first  be  tested  with  the  borax  bead  (§141,  3)  for  the  blue 
bead  of  cobalt  (delicate  and  characteristic  but  obscured  by  the  presence 
of  an  excess  of  nickel  (§132,  7)).  The  sulphides  are  then  dissolved  in  hot 
HC1,  using  a  few  drops  of  HN03  (§141,  1),  and  boiled  to  expel  excess  of 
HN03  :  6CoS  +  12HC1  +  4HN03  ==  6CoCl2  +  3S2  +  4NO  +  8HaO . 
pivide  the  solution  into  three  portions:  To  one  portion  of  the  solution 


190  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES.  §141,  1. 

add  an  excess  (§142,  #)  of  nitroso-  /9-Naphthol,  filter,  and  wash  with  hot 
water  and  then  with  hot  HC1  (§132,  65).  Test  the  red  precipitate  with 
the  borax  bead  for  cobalt.  Render  the  filtrate  ammoniacal,  filter  again 
and  test  this  last  filtrate  with  H2S  for  the  black  precipitate  of  NiS  (§133, 
66  and  e).  To  another  portion  of  the  solution  add  NaHC03  in  excess, 
then  add  H202  ,  warm  and  filter,  a  green  color  to  the  filtrate  indicates 
cobalt  (§132, 10).  The  third  portion  of  the  solution  is  boiled  with  an 
excess  of  NaOH ,  bromine  water  (10,  §§132  and  133)  is  added  and  the  solu- 
tion is  again  boiled.  The  black  precipitate  of  the  higher  hydroxides 
(§141,4)  of  Co  and  Ni  is  thoroughly  washed  with  hot  water  and  then 
treated  on  the  filter  with  hot  solution  of  KI  (§133,  6/),  catching  this  last 
filtrate  in  a  test-tube  containing  CS2  (§141,  6).  Free  iodine  is  evidence  of 
the  presence  of  nickel. 

§141.  Notes. — (1)  HN03  interferes  with  the  nitroso- /3-naphthol  reaction  that 
follows  the  solution  of  the  sulphides  of  Co  and  Ni  ,  hence  an  excess  is  to  be 
avoided.  A  crystal  of  KC1O3  may  be  used  instead  of  HNO3  . 

(2)  If  an  insufficient  amount  of  nitroso- /3-naphthol  has  been  used  a  portion 
of  the  cobalt  maj'  be  in  the  nitrate  and  will  give  the  black  precipitate   for 
nickel.     The    filtrate    must    be    tested    with    the    reagent    to    insure    complete 
removal  of  the  cobalt. 

(3)  Test  with  the  borax  bead  as  follows:  Make  a  small  loop  on  the  end  of  a 
platinum   wire,   dip   this   loop   when   hot  into   powdered    borax,    and    heat   the 
adhering  mass  in  the  flame  until  a  uniform  transparent  glassy  bead  is  obtained. 
Repeat  until  a  bead  the  size  of  a  kernel  of  wheat  has  been  made.     Bring  this 
hot  bead  into  contact  with  the  precipitate  or  solution  to  be  tested  and  fuse 
again  in  the  burner  flame.     Allow  the  bead  to  cool  and  notice  the  appearance. 
A  deep  blue  indicates  cobalt,  obscured,  however,  by  a  large  excess  of  nickel. 

(4)  The  nickel  and  cobalt  may  also  be  oxidized  for  the  KI  test  as  follows: 
Add  five  or  ten  drops  of   bromine  to   the   solution   to  be  tested  in   a   beaker, 
•warm  on  a  water  bath  under  the  hood  until  the  bromine  is  nearly  all  expelled, 
then  add  rapidly  an  excess  of  a  hot  saturated  solution  of  Na2C03  .     The  black 
precipitate  so  obtained  will  filter  rapidly. 

(o)  The  test  for  nickel  by  adding  KI  to  the  mixed  higher  oxides  of  cobalt 
ami  nickel  is  characteristic  of  nickel  and  is  also  a  very  delicate  test.  Fully 
nine-tenths  of  the  cobalt  salts  sold  for  chemically  pure,  show  the  presence  of 
nickel  by  this  test. 

(£)  In  the  reaction  of  nickelic  hydroxide  with  potassium  iodide  some  potas- 
sium iodate  is  formed  and  a  greater  amount  of  free  iodine  will  be  obtained  if 
a  drop  of  hydrochloric  acid  be  added  to  the  filtrate:  KIO3  +  SKI  +  6HC1  = 
3I2  +  6KCf  +  3BVO 

(7)  If  the  sulphides  of  Ni  and  Co  be  digested  with  yellow  ammonium  sul- 
phide, a  portion  of  the  NiS  will  be  dissolved  (§133,  6e)  and  may  be  reprecipi- 
tated  as  a  gray  precipitate  (black  with  free  sulphur)  upon  acidulating  the 
filtrate  with  acetic  acid.  It  is  not  a  delicate  test. 

§142.  Manipulation. — The  solution  of  the  sulphides  of  manganese  and 
zinc  in  cold  dilute  hydrochloric  acid  is  boiled  thoroughly  to  insure  the 
removal  of  the  liydrosulphuric  acid  (§143,  1),  cooled  (§135,  6a),  and  then 
treated  with  an  excess  of  sodium  hydroxide.  The  zinc  forms  the  soluble 
zincate,  Na2Zn02 ,  while  the  manganese  is  precipitated  as  the  hydroxide, 
white,  rapidly  turning  brown  by  oxidation : 

MnCl2  +  2NaOH  =  Mn(OH)2  +  2KC1 

Z&01,  +  4NaOH  =  Na3Zn02  +  gNaCl  -f-  3H.O 


§144.  ANALYSIS  OF  IRON  AND  ZINC  GROUPS.  191 

Filter  and  test  the  filtrate  with  H,S ,  a  white  or  grayish-white  precipitate 
indicates  zinc  (characteristic).  Dissolve  the  well  washed  precipitate  of 
Mn(OH)2  in  nitric  acid  and  boil  with  an  excess  of  lead  peroxide,  adding 
more  nitric  a-cid.  A  violet  color  to  the  nitric  acid  solution  indicates  the 
presence  of  manganese  (very  delicate  and  characteristic) : 

2Mn(OH)2  +  5Pb02  +  10HN03  =  2HMn04  +  5Pb(N03)2  +  GH,0 

§143.  Notes. — 1.  If  the  H2S  is  not  completely  removed  the  Zn  will  be  pre- 
cipitated as  the  sulphide  upon  adding-  the  NaOH  ,  and  will  not  be  separated 
from  the  manganese:  ZnCl2  +  H2S  +  2NaOH  =  ZnS  +  2NaCl  +  2H20  . 

2.  Frequently  the  precipitate  of  zinc  sulphide  is  dark  gray  or  almost  black. 
This  is  usually  due  to  the  presence  of  traces  of  other  sulphides.     If  iron  has  not 
been  all  removed,  through  failure  to  oxidize  completely  with  the  nitric  acid, 
it  may  appear  as  a  precipitate  with  the  manganese,  and  also  as  a  black  precipi- 
tate with  the  zinc  sulphide. 

3.  Small  amounts  of  Co  and  Ni  are  frequently  dissolved  by  the  cold  dilute 
HC1  and  will  appear  with  the  precipitate  o{  Mn(OH)2  .     They  do  not  interfere 
with  the  final  test  for  manganese. 

4.  The  precipitate  of  Mn(OH)2  must  be  washed  to  remove  all  the  chloride, 
as  the  manganese  will  not  be  oxidized  to  permanganic  acid  until  the  chloride 
is  completely  oxidized  to  chlorine. 

•5.  Instead  of  Pb02  ,  red  lead,  Pb304  ,  is  frequently  employed  with  the  nitric 
acid  to  oxidize  the  manganese  to  permanganic  acid: 

2Mn(OH)2  +  5Pb3O4  +  30HN03  =  2HMn04  -f  15Pb(N03)2  +  16H20 

6.  It  is  very  difficult  to  procuie  PbO2  or  Pb304  which  does  not  contain  traces 
of  manganese.     The  student  should  always  boil  the  lead  oxides  with  nitric  acid, 
and  if  a  violet-colored   solution   is  formed,   this   should   be   decanted   and   the 
operation    repeated   until   the    solution   is    perfectly    colorless    after   the    black 
precipitate  of  PbO2  has  subsided.     Then  the  unknown  solution  in  HNO3  may 
be  added  and  the  boiling  repeated  to  test  for  the  manganese. 

7.  The  student  is  not  advised  to  apply  the  permanganate  test  to  the  original 
substances.     All  reducing  agents  interfere,  and  ~M.nO.,  frequently  fails  to  give 
permanganic   acid   when   boiled   with   PbOo    and   HN03    until   after   reduction 
(§134,  6c). 

ANALYSIS  OF  IRON  AND  ZINC  GROUPS  AFTER  PRECIPITATION  BY  AMMONIUM 

SULPHIDE. 

§144.  It  is  preferred  by  some  to  precipitate  the  metals  of  the  third 
and  fourth  groups  together,  by  means  of  ammonium  sulphide;  using 
ammonium  chloride  to  prevent  the  precipitation  of  magnesium  (§189,  55 
and  6a),  and  to  insure  the  complete  precipitation  of  the  aluminum  as  the 
hydroxide  §124,  Go).  In  the  manipulation  for  this  method  of  separation, 
the  H2S  is  not  removed  from  the  second  group-filtrate,  nor  is  nitric  acid 
used  to  oxidize  any  iron  that  may  be  present.  To  the  second  group  filtrate 
(§80),  warmed,  an  excess  of  NH4C1  is  added  (§189,  5c),  then  NH4OH  till 
strongly  alkaline,  and,  paying  no  attention  to  any  precipitate  that  may  be 
formed  (6a,  §§124,  125  and  126),  normal  ammonium  sulphide  is  added  (or 
what  is  equivalent  H2S  is  passed  into  the  alkaline  mixture).  Aluminum 
and  chromium  are  precipitated  as  the  hydroxides,  the  remaining  metals  as 
the  sulphides.  The  following  table  illustrates  a  plan  of  separation  of  the 
ammonium  sulphide  precipitates  of  the  third  and  fourth  group  metals, 
phosphates  being  absent: 


193 


ANALYSIS  OF  IRON  AND   ZING   GROUP. 


144 


o 

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— 

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m  %       g 


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«  ^  'S 


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CO     >-.   C   r^ 

j^.a 


§149.  IRON  AND  ZINC  GROUPS.  103 

§145.  The  presence  of  phosphates  greatly  complicates  the  work  of  the 
analysis  of  the  metals  of  the  third,  fourth,  and  fifth  groups.  The  phos- 
phates of  the  alkali  metals  are  soluble,  those  of  the  other  metals  insoluble 
in  water.  As  the  solutions  for  precipitation  of  first  and  second  group 
metals  are  acid;  phosphates  remain  in  solution  and  do  not  in  any  way 
interfere  with  the  analysis  for  the  metals  of  those  groups;  i.  e.,  silver 
phosphate  in  nitric  acid  solution  is  readily  transposed  by  HC1  ;  copper 
phosphate  in  acid  solution  is  readily  transposed  by  H2S  ;  etc. 

§146.  When  the  filtrate  from  the  second  group  is  rendered  strongly 
ammoniacal  (§128)  the  phosphates  of  all  the  metals  present,  except  those 
of  the  alkalis,  are  precipitated.  Phosphates  of  cobalt,  nickel  and  zinc  are 
redissolved  by  an  excess  of  ammonium  hydroxide.  Freshly  precipitated 
ferric  phosphate  is  transposed  by  the  alkali  hydroxides  (incompletely  in 
the  cold).  The  phosphates  of  Al ,  Cr ,  and  Zn  are  soluble  in  the  fixed 
alkalis,  the  solution  of  chromium  phosphate  is  decomposed  by  boiling, 
precipitating  Cr(OH)3  and  leaving  the  alkali  phosphates  in  solution. 

§147.  In  analysis  a  portion  of  the  filtrate  from  the  second  group  (after 
the  removal  of  the  H2S)  (§128)  should  be  tested  for  phosphoric  acid  with 
ammonium  molybdate  (§75,  6d).  If  phosphates  are^present  the  usual 
methods  of  analysis  for  third,  fourth,  and  fifth  groups  must  be  modified. 
Several  methods  have  been  recommended: 

§148.  First.— To  the  filtrate  from  the  second  group,  H2S ,  being  re- 
moved (§128),  an  excess  of  the  reagent  ammonium  molybdate  is  added, 
the  mixture  set  aside  in  a  warm  place  for  several  hours,  until  the  yellow 
ammonium  phospho-motybdate  has  completely  formed  and  settled 
(§75,  6d).  Filter  and  evaporate  nearly  to  dryness  to  remove  the  nitric  acid. 
Take  up  with  water  and  a  little  hydrochloric  acid  if  necessary  to  obtain  a 
clear  solution,  and  remove  the  excess  of  molybdenum  with  H2S  (§75,  6e). 
From  this  point  proceed  by  the  usual  methods  of  analysis  (§§127,  128 
and  ff.). 

§149.  Second. — Precipitation  of  the  phosphate  as  ferric  phosphate  in 
acetic  acid  solution.  This  method  of  separation  rests  upon  the  fact  that 
the  phosphates  of  the  fourth  group  and  of  the  alkaline  earths  are  soluble, 
and  the  phosphates  of  Al ,  Cr'"  and  Fe'",  insoluble  in  acetic  acid. 

To  the  filtrate  from  the  second  group,  freed  from  H2S  by  boiling  (128), 
and  nearly  neutralized  with  Na2C03 ,  an  excess  of  NaCoH;!02  is  added  and 
then  FeCl3  solution,  drop  by  drop,  as  long  as  a  precipitate  is  formed. 
Care  must  be  taken  to  avoid  an  excess  of  PeCl3 ,  as  the  ferric  phosphate 
is  soluble  in  a  solution  of  ferric  acetate.  As  soon  as  the  phosphate  is  all 
precipitated  the  blood-red  ferric  acetate  is  formed  at  once,  indicating  the 
presence  of  a  sufficient  amount  of  FeCl3 .  The  mixture  should  be  boiled 


194  IRON  AND  ZINC  GROUPS.  §150. 

to  precipitate  the  ferric  acetate  as  basic  ferric  acetate  (§126,  6&)  and  at 
once  filtered. 

Upon  the  addition  of  the  sodium  acetate  the  aluminum  and  chromium 
are  precipitated  as  phosphates,  provided  there  be  sufficient  phosphate 
present  to  combine  with  them;  if  not  the  whole  of  the  phosphate  will  be 
precipitated  and  the  first  drop  of  FeCl3  will  give  a  red  solution  showing 
the  addition  of  that  reagent  to  be  unnecessary. 

By  the  above  method  of  manipulation  any  iron  present  in  the  original 
solution  is  in  the  ferrous  condition  and  does  not  react  to  precipitate  the 
phosphate,  as  ferrous  phosphate  is  soluble  in  acetic  acid.  If  the  iron  has 
been  previously  oxidized  with  nitric  acid  it  will  react  with  the  phosphate 
upon  the  addition  of  the  sodium  acetate;  but  if  there  be  more  iron  present 
than  necessary  to  combine  with  the  phosphate,  the  red  ferric  acetate  solu- 
tion will  be  formed  with  the  excess  of  the  iron  and  render  the  precipita- 
tion of  the  phosphate  incomplete.  In  this  case  the  previous  oxidation  of 
the  iron  is  detrimental. 

If  alkaline  earth  salts  are  present  in  quantity  more  than  sufficient  to 
combine  with  the  phosphoric  acid  radical,  not  all  of  these  metals  will  be 
precipitated  with  the  third  group  metals  upon  the  addition  of  ammonium 
hydroxide.  The  table  (§152)  illustrates  the  separation  of  the  metals  in 
presence  of  the  phosphates  by  the  use  of  FeCl3  in  acetic  acid  solution. 

§150.  Third. — A  method  of  separation  of  the  third  group  metals  with 
phosphates  from  the  remaining  metals  is  based  upon  the  action  of  freshly 
precipitated  barium  carbonate.  Solutions  of  Al,  Cr'",  and  Fe'"  are  pre- 
cipitated as  the  hydroxides  by  digestion  in  the  cold  with  freshly  precipi- 
tated BaC03  (6a/§§124,  125  and  126):  2A1C13  +  3BaC03  +  3H20  = 
2A1(OH)3  -f-  SBaCL  +  3C02 .  Solutions  of  the  chlorides  or  nitrates  of 
the  fourth  group  and  of  the  alkaline  earths  are  not  transposed  by  cold 
digestion  with  BaC03 .  Sulphates  of  the  fourth  group  are  transposed  by 
freshly  precipitated  BaC03  in  the  cold:  CoS04  +  BaC03  =  BaS04  + 
CoC03 ,  etc.;  and  must  not  be  present  in  this  method  of  separation 
(§126,  6a). 

If  an  excess  of  ferric  chloride  be  present  the  phosphates  will  all  be 
precipitated  as  ferric  phosphate  and  the  Al ,  Cr'"  and  excess  of  Fe'"  as 
the  hydroxides  upon  the  digestion  with  BaC03 .  The  table  (§153)  gives 
an  illustration  of  the  use  of  the  BaC03  in  effecting  the  separation. 

It  should  be  observed  that  presence  or  absence  of  FeCl3  or  of  BaC03  in 
the  sample  must  be  fully  determined  before  their  addition  as  reagents. 

§151.  Oxalates  do  not  interfere  with  the  usual  course  of  analysis  of  the 
first  two  groups  of  metals;  with  the  other  metals  oxalates  interfere  very 
much  the  same  as  phosphates.  They,  however,  with  other  interfering 


§151.  IRON  AND  ZINC  GROUPS.  195 

organic  matter,  can  readily  be  removed  by  ignition.  If  the  presence  of 
an  oxalate  has  been  established  (§§188,  6?;  and  227,  8),  the  second  group 
filtrate  should  be  evaporated  to  dryness,  moistened  with  concentrated 
HN03  and  gently  ignited.  The  residue,  dissolved  in  HC1 ,  is  then  ready 
for  the  usual  process  of  analysis.  For  the  analysis  in  presence  of  silicates 
and  borates  the  student  is  referred  to  the  text  under  those  elements 
(§§249,  8  and  221,  8). 


196 


CALCIUM  GROUP  METALS. 


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198  CERIUM—  COLUMBIUM.  §154. 

THE  BARER  METALS  OF  THE  IRON  AND  Zixc  GROUPS. 

Cerium,   Columbium   (Niobium),  Didymiinn,  Erbium,  Gallium,  Glucinum 
(Beryllium),  Indium,  Lanthanum,  Neodymium,  Praseodymium,  Sama- 
rium, Scandium,  Tantalum,  Terbium,  Thallium,  Thorium,  Tita- 
nium, Uranium,  Vanadium,  Ytterbium,  Yttrium,  Zirconium. 

§154.  Cerium.     Ce  =  140.25.    Valence  three  and  four. 

Specific  gravity,  6.628.  Melting  point,  640°  (Cir.  B.  of  S.,  1915).  Cerium  is 
a  comparatively  rare  metal,  never  found  native;  it  is  found  in  many  minerals  in 
Sweden,  especially  in  cerite,  which  is  chiefly  a  silicate  of  Ce  ,  La  ,  Ne  ,  Pr  ,  Al 
and  Fe  ;  also  found  in  a  brick-making  clay  near  Frankfurt,  Germany  (Stro- 
hecker,  J.  pr.,  1886,  (2),  33,  133  and  260).  It  was  first  described  in  1803  by 
Klaproth,  but  in  1839  Mosander  showed  the  supposedly  pure  cerium  oxide  to 
consist  of  oxides  of  at  least  three  metals:  Ce  ,  La  ,  D  (Ne  and  Pr)  (Pogg.,  1842, 
66,  503).  Commercial  "cerium"  consists  of  all  of  these  metals  in  varying  pro- 
portion and  is  known  as  mixed  metal  (misch  metal).  The  metal  is  obtained 
from  the  chloride,  CeCl3  ,  by  electrolysis  or  by  heating  with  sodium.  It  is  a 
steel-gray,  lustrous,  malleable,  ductile  metal;  fairly  stable  in  air  under  ordinary 
conditions.  When  heated  in  air  it  burns  with  incandescence.  The  impure 
commercial  cerium  is  alloyed  with  iron  and  other  heavy  metals  and  used  in 
friction  ignition  devices.  The  iron  alloy  is  known  as  auermetal.  It  burns  in  Cl , 
Br  and  in  vapor  of  I  ,  S  and  P  .  Soluble  in  acids.  Two  oxides  are  known, 
Ce2O3  and  CeO2  ,  forming  two  classes  of  salts,  cerous  and  eerie,  the  latter  being 
less  stable.  Ignition  in  air  or  oxygen  changes  Ce2O3  to  CeO>  .  Ce2O3  is  white 
or  grayish-white,  soluble  in  acids  and  formed  by  igniting  Ce2(CO3)3  ,  Ce_>(C2O4)3 
or  Ce62  in  an  atmosphere  of  hydrogen.  Cerous  salts  are  white  and  form  color- 
less solutions  in  water.  Ceric  oxide,  CeO2  ,  is  yellowish-white,  orange-yellow 
when  hot,  soluble  in  acids  with  difficulty;  the  hydroxide  dissolves  readily. 
Ceric  salts  are  yellow  or  red,  forming  yellow  solutions.  Ceric  hydroxide, 
Ce(OH)4  ,  dissolves  in  HC1  with  evolution  of  chlorine,  forming  colorless  cerous 
chloride.  Sulphurous  acid  decolorizes  solutions  of  eerie  salts,  forming  cerous 
salts.  Fixed  alkali  hydroxides  and  ammonium  sulphide  precipitate,  from  solu- 
tions of  cerous  salts,  the  white  cerous  hydroxide,  turning  yellow  by  absorption 
of  oxygen,  with  formation  of  eerie  hydroxide.  The  precipitate  is  insoluble  in 
excess  of  the  fixed  alkalis  (distinction  from  Al  and  Gl).  The  precipitation  is 
hindered  by  the  presence  of  tartaric  acid  (distinction  from  yttrium).  Ammo- 
nium hydroxide  precipitates  a  basic  salt;  if  H2O2  is  added  before  neutralizing 
a  reddish  brown  precipitate  is  formed  (delicate  test  for  cerium).  Alkali  car- 
bonates precipitate  cerous  carbonate,  soluble  in  excess  of  the  fixed  alkali  car- 
bonates. Oxalic  acid  forms  cerous  oxalate,  white,  from  moderately  acid  solutions, 
soluble  in  hot  (NH4)2C2O4  ,  but  reprecipitated  on  dilution  with  cold  water.  The 
oxalate  is  less  soluble  in  hot  than  in  cold  water.  A  concentrated  solution  of 
K:SO4  forms  the  double  sulphate,  K3Ce(SO4)3  ,  white,  sparingly  soluble  in  water, 
insoluble  in  K  SO,  solution  (distinction  from  Gl).  NajS^O;  does  not  precipitate 
cerium  salts.  BaCO  does  not  precipitate  cerous  salts  in  the  cold,  but  precipitates 
them  completely  on  boiling.  Ceric  salts  are  completely  precipitated  by  BaCO3 
in  the  cold.  Alkali  hypochlorites  precipitate  cerous  salts  as  the  yellow  eerie 
hydroxide.  If  cerous  nitrate  be  boiled  with  PbO2  and  HNO.j  ,  eerie  nitrate,  a 
deep  yellow  solution,  is  formed  (delicate  test  for  cerium).  Cerium  gives  no  absorp- 
tion spectrum,  but  the  spark  spectrum  shows  several  brilliant  lines. 

§155.  Columbium  (Niobium).    Cb  =  93.5.    Valence  five. 

Columbium  usually  occurs  with  tantalum  in  such  minerals  as  columbite  and 
tantalite;  it  is  also  found  in  tantalum  free  minerals  as  euxenite,  pyrochlor,  etc. 
The  metal  is  prepared  by  passing  the  penta-chloride  mixed  with  hydrogen 
repeatedly  through  a  hot  tube.  It  is  a  steel-gray  lustrous  metal,  specific  gravity, 
7.06  at  15.5°  Melting  point,  ±1700  (Cir.  B.  of  S.,  1915).  By  ignition  in  the 


§156.  DIDYMIUM.  199 

air  it  burns  readily  to  the  pentoxide.  Not  attacked  by  chlorine  in  the  cold,  but 
when  warmed  combines  readily,  forming  CbCl5  .  The  metal  is  not  soluble  in 
hydrochloric,  nitric  or  nitrohydrochloric  acid,  but  is  readily  soluble  in  hot  con- 
centrated sulphuric  acid,  forming  a  colorless  solution  (Roscoe,  C.  N.,  1878,  37, 
25).  It  forms  several  oxides,  CbO  ,  CbO..  and  Cb.O5  .  Columbic  acid  (anhy- 
dride) Cb_O5,  is  a  white  powder,  yellow  when  hot  (distinction  from  tantalum); 
it  is  obtained  by  ignition  of  the  lower  oxides,  or  by  decomposition  of  solutions 
of  the  salts  by  water  or  alkalis  and  igniting.  CbOj  ,  black,  is  prepared  by 
strongly  igniting  Cb>O5  in  a  current  of  hydrogen.  Cb^Os ,  not  too  strongly 
ignited,  is  soluble  in  acids,  from  which  solutions  NH4OH  and  (NH4)2S  pre- 
cipitate columbic  acid  containing  some  ammonia.  By  mixing  CbzOs  with  char- 
coal and  heating  in  a  current  of  chlorine,  a  mixture  of  CbOCl:i  and  CbCls  is 
obtained.  CbCl5  is  a  yellow  crystalline  solid  (needles),  melting  at  194°  and 
distilling  at  240.5°  (Deville  and  Troost,  C.  r.,  1867,  64,  294).  Upon  treating 
the  chloride  with  water,  it  is  partially  decomposed  to  columbic  acid,  a  large 
portion  remaining  in  solution  and  not  precipitated  by  H2SO4  (distinction  from 
tantalum).  Cb^O^  not  previously  ignited  dissolves  in  HF  ;  which  solution, 
when  mixed  with  KF  ,  the  HF  being  in  excess,  gives  a  double  fluoride,  2KF.CbF5; 
if  the  HF  be  not  in  excess,  a  double  oxy-fluoride  is  obtained,  2KF.CbOF3  (Kruess 
and  Nilson,  B.,  1887,  20,  1676).  The  potassium  columbium  fluoride  is  much 
more  soluble  than  either  the  corresponding  titanium  or  tantalum  compounds. 
Fusion  of  columbic  acid  with  the  alkalis  gives  the  columbates,  the  potassium 
salt  being  quite  soluble  in  water  and  in  potassium  hydroxide;  the  s  dium  salt 
is  only  soluble  in  water  after  removal  of  the  excess  of  the  sodium  hydroxide. 
From  a  solution  of  potassium  columbate,  sodium  hydroxide  precipitates,  almost 
completely,  sodium  columbate.  Carbon  dioxide  precipitates  columbic  acid  from 
solutions  of  columbates.  Soluble  salts  of  Ba  ,  Ca  and  Mg  form  white  bulky 
precipitates  with  a  solution  of  potassium  columbate.  AgNO3  gives  a  yellowish-white 
precipitate,  CuSO4  a  green  precipitate.  Cb2O5  in  presence  of  HC1  or  H  SO,, 
gives  a  blue  to  brown  color  with  Sn  or  Zn,  due  to  partial  reduction  of  the  Cb  (dis- 
tinction from  tantalum).  Fused  with  sodium  meta-phosphate,  columbic  acid 
gives  in  the  inner  flame  a  violet  to  blue  bead;  a  red  bead  by  addition  of  FeSOj. 

(  Neodymium.    Nd  =  144.3.    Valence  three. 
§156.  Didymmm  =  i  _  _  n  _      ,T  , 

(Praseodymium.    Pr=140.9.     Valence  three. 

Specific  gravity,  6.544.  Melting  points,  Neodymium,  840?;  praseodymium, 
940?  (Cir.  B.  of  S.,  1915).  Present  in  cerite  in  Sweden  and  in  monazite  sand 
from  Brazil.  Didymium  was  reported  about  1840  by  Mosander,  having  been 
separated  from  cerium  and  lanthanum.  In  1885  Welsbach  (M.,  1885,  6,  477) 
separated  didymium  salts  into  two  distinct  salts,  neodymium  and  praseody- 
mium. By  the  absorption  spectrum  bands  other  chemists  are  of  the  opinion 
that  the  so-called  didymium  consists  of  a  group  of  elements,  nine  or  more  (Kruess 
and  Nilson,  B.,  1887,  20,  2166;  Kruess,  A.,  1892,  265,.  1).  Concerning  the 
separation  of  didymium  compounds,  see  Dennis  and  Chamot  (/.  Am.  Soc.,  1897, 
19,  799).  By  repeated  fractionation  of  the  nitrate  (several  thousand  times) 
Welsbach  obtained  a  pale  green  salt  and  a  rose-colored  salt,  which  gave  dif- 
ferent spectra,  but  which,  united,  gave  the  spectrum  of  didymium.  Didymium 
oxide  absorbs  water  to  form  the  hydroxide,  which  absorbs  CO2  from  the  air, 
but  does  not  react  alkaline  to  litmus.  The  salts  are  soluble  in  water  to  a  reddsih 
solution.  The  saturated  sulphate  solution  does  not  deposit  crystals  until  heated 
to  boiling;  while  lanthanum  sulphate  precipitates  from  the  saturated  solution 
at  30°.  Fixed  alkalis  precipitate  the  hydroxide:  NH4OH  ,  a  basic  salt;  insoluble 
in  excess  of  the  reagents.  Alkali  carbonates  form  a  bulky  precipitate,  insoluble 
in  excess  of  the  reagent,  barium  carbonate  precipitates  slowly  but  completely. 
Precipitation  by  alkalis  is  prevented  by  tartaric  acid.  Oxalic  acid  precipitates 
didymium  salts  completely,  soluble  with  difficulty  in  HC1 .  The  double  potas- 
sium sulphate  forms  much  more  slowly  and  less  completely  than  with  cerium. 
The  salts  give  a  distinct  and  characteristic  absorption  spectrum.  Consult  Jones, 
(Am.,  1898,  20,  345),  Schele  (Z.  anorg.,  1898,  17,  319),  Boudard  (C.  r.,  1898, 
126,  900),  Demarcay  (C.  r.,  1898,  126,  1039),  and  Brauner  (C.  N.,  1898,  77,  161). 


200  ERBIUM— GALLIUM— GLUCINUM.  §157. 

§157.  Erbium.     Er  =  167.7.     Valence  three. 

Erbium  has  been  prepared  in  the  form  of  a  dark  gray  powder.  Specific  gravity, 
4.77  at  15°.  (Meyer,  Monatsch.,  20,  793,  1899).  As  oxide  or  earth  it  is  de- 
scribed by  Cleve  (C.  r.,  1880,  91,  381)  as  that  yttrium  earth  the  most  beautiful 
rose  colored.  It  forms  a  characteristic  absorption  spectrum,  and  a  spark  spec- 
trum with  sharp  lines  in  tfie  orange  and  green.  This  earth  has  not  been  thor- 
oughly studied  and  quite  probably  consists  of  the  oxides  of  several  metals  (Bois- 
baudran,  C.  r.,  1886,  102,  1003;  Soret,  C.  r.,  1880,  91,  378;  Crookes,  C.  N.,  1886, 
64,  13).  The  oxide  gives  upon  ignition  an  intense  green  light;  it  is  not  fusible 
or  volatile. 

§158.  Gallium.    Ga=69.9.    Valence  three. 

Specific  gravity,  the  solid,  at  23°  to  24.5°,  5.935  to  5.956;  the  melted,  at  24.7°, 
6.069.  Melting  point,  30.15°;  frequently  may  be  cooled  to  0°  without  again  be- 
coming solid.  It  is  a  grayish-white  metal,  crystallizing  in  octahaedra  or  in 
broad  plates.  It  is  quite  brittle  and  gives  a  bluish-gray  mark  on  paper.  It- 
gives  a  very  weak  and  fugitive  flame  spectrum;  the  spark  spectrum  shows  two 
beautiful  violet  lines.  When  heated  in  the  air  or  in  oxygen  it  is  but  slightly 
oxidized;  does  not  vaporize  at  a  white  heat;  soluble  in  acids  and  alkalis; 
attacked  by  the  halogens  (with  iodine  only  upon  warming).  In  the  Periodic 
System  it  is  the  Ekaaluminum  of  Mendelejeff,  who  described  the  general  prop- 
erties before  the  metal  was  discovered  (C.  r.,  1875,  81,  969).  It  occurs  in  zinc 
blende  (black)  from  Bensberg  on  the  Rhine;  in  brown  blende  from  the 
Pyrenees;  and  in  some  American  zinc  blendes  (Cornwall,  Ch.  Z.,  1880,  4,  443). 
It  is  prepared  by  electrolysis  after  previous  purification  of  the  ore  by  chemical 
methods.  4300  kilos  of  the  Bensberg  ore  gave  55  kilos  of  pure  gallium  (Bois- 
baudran  and  Jungfleisch,  C.  r.,  1878,  86,  475).  The  oxide,  Ga,O3  ,  is  a  white 
powder  obtained  by  igniting  the  nitrate.  After  strong  ignition  it  is 
insoluble  in  acids  or  alkalis.  It  is  easily  attacked  on  fusion  with  KOH 
or  KHSO4  .  The  alkalis  and  the  alkali  carbonates  precipitate  the  salts 
as  the  hydroxide,  perceptibly  soluble  in  fixed  alkali  carbonates,  more  easily 
in  ammonium  hydroxide  and  in  ammonium  carbonate,  and  very  readily  in 
the  fixed  alkalis.  Tartrates  hinder  the  precipitation  of  the  hydroxide.  The 
salts  of  gallium  are  colorless  and  for  the  most  part  soluble  in  water.  The 
neutral  solutions  upon  warming  precipitate  a  basic  salt,  dissolving  again  upon 
cooling.  Excess  of  zinc  forms  a  basic  zinc  salt  which  precipitates  the  gallium 
as  oxide  or  basic  salt.  BaC03  precipitates  gallium  salts  in  the  cold.  K4re(CN)8 
gives  a  precipitate,  insoluble  in  ITC1 ,  noticeable  in  very  dilute  solutions 
(1-175,000).  H,S  does  not  precipitate  gallium  salts  from  solutions  acid  with 
mineral  acids;  from  the  acetate  or  in  presence  of  ammonium  acetate  the  n'lnte 
sulphide,  Ga2S8  ,  is  precipitated;  (NH4)2S  precipitates  the  sulphide.  Gallium 
chloride,  GaCL,  ,  is  a  colorless  salt,  melting  at  75°  and  volatilizing  at  215°  to 
220°.  The  vapor  density  indicates  the  molecule  to  be  Ga2Cl6 ,  which  decomposes 
to  GaCl3  at  about  400°  (Friedel  and  Kraft,  C.  r.,  1888,  107,  306).  Upon  evaporat- 
ing a  solution  of  the  chloride  on  a  water  bath  the  salt  is  perceptibly  volatil- 
ized, not  so  if  HzSOj  be  present.  Gallium  sulphate  forms  with  ammonium 
sulphate  an  alum.  For  separation  from  other  metals,  see  Boisbaudran,  C.  r., 
1882,  95,  410,  503,  1192,  1332. 


§159.  Gluciimm  (Beryllium).    Gl  =  9.1 .    Valence  two. 

Specific  gravity,  1.85  (Humpidge,  Proc.  Roy.  Soc.,  1871,  39,  1).  Melting  point, 
1350°  ?  (Cir.  B.  of  S.,  1915).  It  is  a  white  malleable  metal,  obtainable  in  hexagonal 
crystals  (Nilson  and  Pettersson,  B.,  1878,  11,  381  and  906).  It  was  first  dis- 
covered in  1797  by  Vauquelin  from  beryl.  The  powdered  metal  takes  fire  when 
heated  in  air,  burning  with  great  brilliancy.  It  dissolves  readily  in  dilute  acids 
and  also  in  alkalis  with  evolution  of  hydrogen.  It  does  not  decompose  steam 
even  at  a  red  heat.  It  is  a  strongly  positive  element,  in  general  properties 
between  aluminum  and  the  alkaline  earths;  as  lithium  is  between  the  alka- 


§160.  INDIUM.  201 

line  earths  and  the  alkali  metals.  It  should  be  classed  with  the  alkaline  earths. 
It  is  found  in  chrysoberyl,  G1(A1O2)2  ,  in  phenakite,  GLSiO,  ,  and  in  some  other 
silicates.  It  is  prepared  by  heating  the  chloride,  G1C12  ,  with  Na  in  a  closed  iron 
crucible  (Nilson  and  Pettersson,  /.  c.);  or  by  heating  the  oxide,  G1O,  with  Mg 
(Winkler,  B.,  1890,  23,  120).  The  oxide,  G1O  ,  is  obtained  by  igniting  the 
hydroxide.  It  is  a  white  infusible  powder,  soluble  in  acids  and  in  fixed  alkalis. 
The  hydroxide  is  prepared  by  precipitating  the  salts  with  NH4OH  ,  soluble 
in  the  fixed  alkalis  and  in  ammonium  carbonate,  concentrated;  precipitated 
on  dilution  and  boiling  (distinction  and  separation  from  Al).  The  metal  is 
soluble  in  acids  except  that  when  in  the  compact  form  it  is  scarcely  attacked 
by  HNOs  .  The  hydroxide  is  soluble  on  continued  boiling  writh  NH4C1  ,  form- 
ing G1C12  .  The  more  common  salts  of  glucinum  are  soluble  in  water  to  a 
solution  having  a  sweetish  taste.  The  carbonate  and  phosphate  are  insoluble, 
the  oxalate  and  sulphate  soluble,  the  existence  of  a  sulphide  is  doubtful.  Solu- 
tions of  glucinum  salts  are  precipitated  by  the  alkalis,  the  precipitate  being 
soluble  in  excess  of  the  fixed  alkalis.  The  alkali  carbonates  precipitate  the 
carbonate,  soluble  in  concentrated  ammonium  carbonate,  reprecipitated  on 
diluting,  boiling  and  adding  an  excess  of  NH4OH  (Joy,  Am.  S.,  1863,  (2),  36,  83). 
The  salts  are  not  precipitated  by  H2S  ,  but  are  precipitated  by  (NH4)2S  as  the 
hydroxide.  BaCO3  does  not  precipitate  Gl  salts  in  the  cold,  but  precipitates 
them  upon  boiling.  G1C12  melts  at  about  600°  and  sublimes  at  a  white  heat, 
forming  white  needles.  The  oxide  has  not  been  melted  or  sublimed.  Gl  usually 
occurs  as  a  silicate  with  aluminum.  The  mass  is  fused  with  alkali  carbonate, 
acidified  with  HC1  and  the  Al  and  Gl  chlorides  filtered  from  the  SiO2  .  An  excess 
of  ammonium  carbonate  precipitates  both  metals,  but  redissolves  the 
Gl  .  After  repeating  this  separation  several  times  pure  glucinum  hydroxide, 
G1(OH)2  ,  is  obtained  upon  boiling  off  the  ammonia.  The  hydroxide  thus  obtained 
is  ignited  and  weighed  as  the  oxide. 

§160.  Indium,     In  •=  114.8.     Valence  three. 

Specific  gravity,  7.11  to  7.28  at  20.4°.  Melting  point,  155°  (Cir.  B.  of  S.,  1915). 
Indium  was  discovered  in  Freiberg  zinc  blende  by  Reich  and  Richter  (/.  pr., 
1863,  89,  441;  90,  175;  1864,  93,  480),  by  use  of  the  spectroscope.  It  is  found 
chiefly  as  sulphide,  never  native,  in  the  Freiberg  blende  to  the  extent  of  about 
0.1  per  cent.  It  is  found  in  a  few  other  places,  but  in  much  smaller  amounts 
(Boettger.  /.  pr.,  1866,  98,  26).  In  the  preparation  of  indium  the  Freiberg 
zinc  is  dissolved  in  HC1  or  H2SO4  ,  leaving  an  excess  of  the  zinc.  When  no 
more  hydrogen  is  evolved,  the  mass  is  digested  for  a  day  or  more  with  the 
excess  of  Zn  ,  whereby  the  indium  is  obtained  as  a  precipitate  with  Pb  ,  Cu  , 
Cd ,  Sn ,  As  ,  Fe  and  Zn  .  This  precipitate  is  dissolved  in  nitric  acid  and 
evaporated  with  sulphuric  acid;  then  taken  up  with  water,  separating  from 
lead.  The  solution  is  precipitated  with  NH4OH ,  which  precipitates  the  In 
and  Fe  ;  this  precipitate  is  dissolved  in  HC1  and  boiled  for  some  time  with 
NaHSO;,  .  The  indium  sulphite  is  obtained  as  a  fine  crystalline  powder,  which 
is  treated  with  HNO3  and  H2SO4  ,  forming  indium  sulphate,  from  which  the 
metal  is  precipitated  by  zinc  (Bayer,  A.,  1871,  158,  372;  Boettger,  /.  pr., 
1869,  107,  39;  Winkler,  J.  pr.,  1867,  102,  276).  Indium  is  a  grayish-white 
metal,  very  soft,  makes  a  good  mark  on  paper,  is  ductile,  easily  fusible,  less 
volatile  than  Zn  or  Cd  .  It  is  less  electro-positive  than  Zn  or  Cd  and  hence 
it  is  precipitated  from  its  solutions  by  both  these  elements.  In  the  air  pr  in 
water  it  is  rather  more  stable  than  zinc.  Heated  in  the  air  it  burns  with  a 
violet  flame  and  brown  smoke,  forming  the  oxide,  In2O3  .  Indium  does  not 
decompose  water  at  100°.  At  a  red  heat  it  combines  with  sulphur  and  the 
halogens.  By  ignition  with  charcoal  or  in  a  current  of  hydrogen  it  is  reduced 
to  the  metal  from  its  compounds.  It  is  soluble  in  HC1  and  H2SO4  ,  evolving 
H  ;  in  HNOS  ,  evolving  NO  .  In  the  reactions  of  its  salts  indium  deports 
itself  quite  similar  to  Fe"'  and  Al .  Its  most  characteristic  property  is  its 
spectrum;  two  lines,  an  indium  a,  intense  blue,  and  an  indium  0,  less  intense 
violet  (Schroetter,  J.  pr.,  1865,  95,  441).  In2O  is  brown  when  hot,  light  yellow 
when  cold,  slowly  soluble  in  cold  acids,  rapidly  when  heated.  Indium  salts 
are  precipitated  by  the  alkalis  as  In(OH)3 ,  soluble  in  excess  of  the  fixed  alkalis, 


202  LANTHANUM— SCANDIUM.  §161. 

reprecipitated  by  boiling  or  treating  with  NH4C1 .  Tartrates  prevent  the 
precipitation  by  alkalis.  Alkali  carbonates  precipitate  the  indium  carbonate, 
soluble  in  ammonium  carbonate,  but  reprecipitated  on  boiling.  BaCO:i  pre- 
cipitates the  indium  completely  as  a  basi  salt  (separation  from  Co  ,  Ni  ,  Mn  , 
Zn  and  Fe").  Phosphates  form  white  precipitates  from  neutral  solutions.  H2S 
precipitates  from  neutral  solutions,  or  solutions  acid  with  acetic  acid,  yellow 
indium  sulphide.  In  alkaline  solutions  H2S  ,  or  in  neutral  solutions  (NH4)2S  , 
forms  a  white  precipitate  containing  In2S3 .  Yellow  In2S3  boiled  with  (NH4)2SX 
becomes  white  and  is  partly  dissolved.  Upon  cooling  the  solution  a  bulky  white 
precipitate  separates  out.  K4Fe(CN)6  gives  a  white  precipitate;  K2CrO4  gives 
a  yellow  precipitate;  K2Cr2O7  ,  K3Fe(CN)6  and  KCNS  do  not  form  precipitates. 

§161.  Lanthanum,    La  =  139.0.    Valence  three. 

Specific  gravity,  6.163.  Melting  point,  810°  ?  (Cir.  B.  of  S.,  1915).  In  general 
appearance  and  properties  very  similar  to  Ce  .  It  is  prepared  almost  exclusively 
from  cerite.  By  treating  the  mineral  with  an  insufficient  quantity  of  HNO3  , 
a  solution  rich  in  La  may  be  obtained.  The  cerium  is  precipitated  from  the 
solution  by  alkali  hypochlorite.  The  filtrate  is  converted  into  the  sulphate  and 
separated  from  Ne  and  Pr  sulphates  by  fractional  crystallization,  the  latter 
being  more  soluble  (Holzman,  J.  pr.,  1858,  75,  346).  Fractional  precipitation 
with  NH4OH  is  also  used  to  separate  La  from  Ne  and  Pr  ,  the  latter  precipitat- 
ing first  (Cleve,  Bl,  1874,  21,  196;  1883,  39,  287).  The  metal  is  prepared  from 
the  chloride,  LaCl3 ,  by  electrolysis  or  by  ignition  with  potassium.  The  igni- 
tion point  of  La  is  higher  than  that  of  Ce  ;  it  is  also  not  so  readily  attacked 
by  HNO3  .  In  cold  water  La  is  slowly  attacked,  but  in  hot  water  the  action 
is  violent  (Winkler,  B.,  1890,  23,  787).  With  aluminium,  lanthanum  forms  a 
crystalline  white  alloy  which  is  stable  in  air  and  insoluble  in  nitric  acid  (Muth- 
man  and  Beck,  A.,  46,  331,  1904.)  The  oxide,  La^Os  ,  is  a  white  powder,  readily 
soluble  in  acids;  with  water  it  forms  the  hydroxide,  La(OH)3  ,  which  reacts 
alkaline  towards  litmus  and  absorbs  CO2  from  the  air.  La(OH)3  is  soluble  in 
a  solution  of  NH4C1  (similar  to  Mg(OH)2).  The  salts  are  colorless.  K2SO4 
and  H2C2O4  form  precipitates  with  lanthanum  salts  as  with  cerium  salts.  Fixed 
alkalis  precipitate  lanthanum  salts  as  La(OH)3  ,  white,  insoluble  in  excess  of 
the  reagent  and  not  changing  color  on  exposure  to  the  air  (distinction  from  Ce). 
Alkali  carbonates  precipitate  La2(CO3)3  ,  insoluble  in  excess.  BaCO3  precip- 
itates the  salts  completely  in  the  cold.  NH4OH  precipitates  basic  salts.  H2S 
forms  no  precipitate;  (NH4)2S  precipitates  the  hydroxide.  Lanthanum  gives  a  num- 
ber of  characteristic  lines  in  the  spark  spectrum  (Bettendorf,  A.,  1889,  256,  159). 

§162.  Neodymium.    Nd  =  144.3.     See  Didymium  (§156). 
§163.  Praseodymium.    Pr  =  140.9.     See  Didymium  (§156). 

§164.  Samarium.     Sa  =  150.4.     Valence  three. 

Samarium  was  found  in  1879  by  Boisbaudran  from  didymium  earths  by  its 
peculiar  spectrum  (C.  r.,  1879,  88,  323).  According  to  Crookes,.  (C.  r.  1886,  102, 
1464),  it  consists  of  at  least  two  elements  and  is  found  in  all  yttrium  earths. 
Its  salts  are  light  yellow,  giving  an  absorption  spectrum  of  six  bands  (Kruess, 
B.,  1887,  20,  2144).  In  its  chemical  properties  it  is  more  similar  to  Nd  and  Pr 
than  to  Y  .  It  is  separated  from  Nd  and  Pr  by  the  fractional  precipitation  of 
the  hydroxide,  basic  nitrate,  oxalate  and  sulphate;  which  separate  before  the  corre- 
sponding Nd  and  Pr  compounds.  Melting  point,  1300°-1400°  (Cir.  B.  of  S.,  1915). 

§165.     Scandium.     Sc  =  44.1.    Valence  three. 

It  is  found  in  euxenite  and  gadolinite  with  yttrium.  Its  name  comes  from 
Scandinavia,  where  it  was  first  found.  It  is  separated  from  ytterbium,  with 
which  it  is  always  closely  associated,  by  heating  the  nitrates;  the  basic  scan- 
dium nitrate  being  precipitated  before  the  ytterbium  basic  nitrate,  or  by 
precipitating  as  the  double  potassium  sulphate,  the  corresponding  ytterbium 


§167.  TANTALUM— TERBIUM.  203 

salt  remaining1  in  solution.  The  oxide,  Sc,03  ,  is  a  white  flocculent  infusible 
powder,  readily  soluble  in  warm  acids.  The  solutions  of  the  salts  show  m> 
absorption  bands  in  the  spectrum.  The  spark  spectrum  of  the  chloride  gives 
over  100  bright  lines  (Thalen,  C.  r.,  1880,  91,  45).  Solutions  of  the  salts  taste 
sweet  and  have  an  astringent  action.  The  alkalis  precipitate  the  hydroxide, 
a  white  bulky  precipitate,  insoluble  in  excess  of  the  precipitant.  Tartrates 
hinder  the  precipitation  in  the  cold,  but  not  upon  heating-.  NaoCO;;  gives  a 
bulky  white  precipitate,  soluble  in  excess  of  the  reagent.  H,S  is  without 
action,  but  (NH4)2S  precipitates  the  ]i]/<Jro.ride.  K,S04  precipitates  the  double 
scandium  sulphate,  3K,S04.Sc2(S04)3 ,  soluble  in  water  but  not  in  a  saturated 
K,S04  solution. 

§166.  Tantalum.    Ta  =  181.5.     Valence  five. 

Tantalum  occurs  in  tantalite  and  columbite,  silicates,  nearly  always  accom- 
panied by  columbium.  It  is  prepared  by  heating  the  tantalum  alkali  fluoride 
with  K  or  Na  in  a  well-covered  crucible  (Rose,  Pogg.,  1856,  99,  65).  The  electric 
furnace  is  now  used  in  its  preparation.  It  is  a  white  metal,  somewhat  less  bright 
than  platinum  and  usually  showing  a  bluish  tarnish  from  superficial  oxidation. 
Specific  gravity,  16.8  (von  Bolton).  Melting  point,  2850°  (Cir.  B.  of  S.,  1915). 
Wrought  tantalum  is  about  as  hard  as  soft  steel,  but  possesses  a  much  greater 
tensile  strength.  When  heated  it  becomes  softer  and  can  be  hammered,  rolled 
and  drawn.  Tantalum  has  become  very  important  commercially  on  account 
of  its  high  melting  point,  its  property  of  hardening  and  improving  steel,  even 
when  added  in  extremely  small  quantities,  and  its  great  acid-resisting  prop- 
erties. At  a  dull  red  heat  tantalum  absorbs  large  quantities  of  hydrogen  and 
nitrogen,  becoming  brittle  and  generally  changing  its  properties.  The  occluded 
gases  may  be  removed  by  fusion  in  vacuo  in  the  electric  furnace.  Heated  in  the 
air  it  burns  with  incandescence  to  form  Ta^Os .  It  is  insoluble  in  acid,  including 
boiling  aqua  regia,  except  HF  ,  in  which  it  dissolves  with  evolution  of  H  .  Upon 
ignition  in  a  current  of  chlorine,  TaCl5  ,  volatile,  is  formed.  Solution  of  alkalis 
has  no  action;  upon  fusion  with  the  fixed  alkalis  an  alkali  tantalate  is  formed. 
Ta2O5  is  a  white  infusible  powder,  specific  gravity,  8.01  (Marignac,  A.  Ch.,  1866, 
(4),  9,  254).  The  oxide  fused  with  fixed  alkalis  gives  also  an  alkali  tantalate, 
M'TaO3  .  When  KOH  is  used,  the  fused  mass  is  soluble  in  water.  When  NaOH 
is  used,  water  removes  the  excess  of  alkali,  leaving  the  NaTaO3  as  a  white 
residue,  which  dissolves  in  pure  water,  but  not  in  NaOH  solution.  Tantalum 
chloride  is  a  yellow  solid,  melting  at  211.3°  and  boiling  at  241.6°,  with  753 
mm.  atmospheric  pressure  (Deville  and  Troost,  C.  r.,  1867,  64,  294).  It  is  com- 
pletely decomposed  by  water,  forming  the  hydrated  acid,  2HTaO3.H2O  = 
H4Ta2O7 .  The  freshly  precipitated  acid  is  soluble  in  acids  and  reprecipitated 
by  NH4OH  .  The  acid  is  readily  soluble  in  HF  ,  which  solution  with  KF  forms 
a  characteristic  double  salt,  2KF.TaF5  ,  crystallizing  in  fine  needles,  insoluble  in 
water  slightly  acidulated  with  HF  (distinction  and  separation  from  colum- 
bium). A  solution  of  alkali  tantalate  gives  with  HC1  a  precipitate  of  tantalic 
acid,  soluble  in  excess  of  the  HC1  .  From  this  solution  NH4OH  or  (NH4)2S 
precipitates  tantalic  acid;  H2SO4  precipitates  tantalic  sulphate.  Tartaric  acid 
prevents  the  precipitation  -with  NH4OH  and  (NH4)2S  .  A  solution  of  tantalic 
acid  gives  no  coloration  with  zinc  (distinction  from  Cb).  Solutions  of  alkali 
tantalates  form  tantalic  acid  with  CO2  .  The  acid  fused  with  sodium  metaphos- 
phate  gives  a  colorless  bead  (distinction  from  SiO2),  which  does  not  become  blood- 
red  upon  adding  FeSO4  and  heating  in  the  inner  flame  (distinction  from  titanium). 

§167.  Terbium.     Tb  =  159.2.    Valence  three. 

The  terbium  compounds  are  very  similar  to  the  yttrium  compounds.  The 
salts  are  colorless  and  give  no  absorption  spectrum.  The  double  potassium 
terbium  sulphate  has  about  the  same  solubilities  as  the  corresponding  cerium 
compound,  and  so  the  terbium  is  frequently  precipitated  with  cerium  com- 
pounds. Terbia,  Tr2O3 ,  is  the  darkest  colored  of  the  yttrium  earths,  soluble 


204  THALLIUM— THORIUM.  £168 

in  acids  and  sets  NH3  free  from  ammonium  salts.  The  hydroxide  is  a 
gelatinous  precipitate  which  absorbs  CO2  from  the  air.  It  is  quite  probable 
that  terbia  is  a  mixture  of  rare  earths  (Boisbaudran,  C.  r.,  1886,  102,  153,  395, 
483  and  899). 


§168.  Thallium.     Tl  =  204.00.     Valence  one  and  three. 

Thallium  was  discovered  by  Crookes  by  means  of  the  spectroscope  in  1861. 
in  selenium  residues  of  the  H2SOt  factory  at  Tilkerode  in  the  Hartz  Mountains, 
Germany  (C.  N.,  1861,  3,  193,  303;  1863,  7,  290;  1863,  8,  159,  195,  219,  231,  243, 
255  and  279).  It  is  found  widely  distributed  in  many  varieties  of  iron  and 
copper  pyrites,  but  in  large  proportions  it  is  only  found  in  Crookesite  in 
Sweden.  This  mineral  contains  as  high  as  18.55  per  cent  Tl  (Nordenskjoeld, 
A.,  1867,  144,  127).  It  is  prepared  by  reduction  from  its  solutions  with  Zn  or 
Al;  by  electrolysis;  by  precipitation 'with  KI  ,  and  then  reduction  by  Zn  or  Al 
or  by  electrolysis.  Specific  gravity,  11.85  (Petrenko,  Z.  anorg.,  50,  133,  1906)' 
Melting  pointy  302°  (Cir.  B.  of  S.,  1915).  It  is  a  bluish-white  metal,  softer 
than  lead,  malleable  and  ductile;  tarnishes  rapidly  in  the  air;  may  be  pre- 
served under  water,  which  it  does  not  decompose  below  a  red  heat ;  soluble 
in  H2S04  and  HNO3  ,  in  HC1  with  great  difficulty;  combines  directly  with 
Cl ,  Br ,  I  ,  P  ,  S  ,  Se ,  and  precipitates  from  their  solutions  Cu. ,  Ag  ,  Hg , 
Au  and  Pb  in  the  metallic  state.  As  a  monad  its  compounds  are  stable,  and 
not  easily  oxidized;  as  a  triad  it  is  easity  reduced  to  the  univalent  condition. 
Thallious  oxide,  TLO  ,  is  black;  on  contact  with  water  it  forms  an  hydroxide, 
T1OH  ,  freely  soluble  in  water  and  in  alcohol,  to  colorless  solutions.  The  car- 
bonate is  soluble  in  about  20  parts  of  water;  the  sulphate  and  phosphate  are 
soluble;  the  chloride  very  sparingly  soluble;  the  iodide  insoluble  in  water. 
Hydrochloric  acid  precipitates,  from  solutions  not  very  dilute,  thalliuiix 
chloride,  T1C1 ,  white,  and  unalterable  in  the  air.  As  a  silver-group  precipitate, 
thallious  chloride  dissolves  enough  in  hot  water  to  give  the  light  yellow  pre- 
cipitate of  iodide,  Til,  on  adding  a  drop  of  potassium  iodide  solution,  the 
precipitate  being  slightly  soluble  in  excess  of  the  reagent.  ELS  precipitates 
the  acetate,  but  not  the  acidified  solutions  of  its  other  salts.  (NH4)oS  pre- 
cipitates T13S ,  which,  on  exposing  to  the  air,  soon  oxidizes  to  sulphate. 
Ferrocyanides  give  a  yellow  precipitate,  Tl4Pe(CN)6;  phosphomolybdic  acid  a 
yellow  precipitate;  and  potassium  permanganate  a  red-brown  precipitate,  con- 
sisting in  part  of  TL03  .  Chrornates  precipitate  yellow  normal  chromate;  and 
platinic  chloride,  pale  orange,  thallious  platinie  chloride,  TLPtClc .  Thallium 
compounds  readily  impart  an  intense  green  color  to  the  flame,  and  one  emerald- 
green  line  to  the  spectrum  (the  most  delicate  test).  The  flame-color  and 
spectrum,  from  small  quantities,  are  somewhat  evanescent,  owing  to  rapid 
vaporization.  Thallic  oxide,  T12O3  ,  dark  violet,  is  insoluble  in  water;  the 
hydroxide,  an  oxyhydroxide,  TIO(OH),  is  brown  and  gelatinous.  This  hydrox- 
ide is  precipitated  from  thallic  salts  by  the  caustic  alkalis,  and  not  dissolved 
be  excess.  Chlorides  and  bromides  do  not  precipitate  thallic  solutions;  iodides 
precipitate  Til  Avith  I.  Sulphides  and  ILS  precipitate  thallious  sulphide,  with 
sulphur.  Thallic  oxide,  suspended  in  solution  of  potassium  hydroxide,  and 
treated  with  chlorine,  develops  an  intense  violet-red  color.  Thallic  chloride 
and  sulphate  are  reduced  to  thallious  salts  by  boiling  their  water  solutions. 

§169.  Thorium.     Th  =  232.4.     Valence  four. 

Thorium  is  a  rare  element  found  in  thorite  (a  silicate),  orangite,  monazite, 
and  some,  other  minerals.  It  was  described  by  Berzelius  in  1828  (Pogg.,  1829, 
16,  385),  who  also  prepared  the  metal  by  reduction  of  the  potassium  thorium 
fluoride  with  potassium.  The  metal  is  a  gray  powder;  specific  gravity,  11.000; 
melting  point,  >1700°,  <1755°  (Cir.  B.  of  S.,  1915);  stable  in  air  at  ordi- 
nary temperature,  but  igniting  when  heated;  attacked  by  vapors  of  Cl  , 
Br  ,  I  and  S  .  Sparingly  soluble  in  dilute  acids,  easily  soluble  in  concentrated 
acids;  insoluble  in  the  alkalis  (Nilson,  B.,  1882,  15,  2519  and  2537;  Kruess 


§170.  TITANIUM.  205 

and  Nilson,  B.,  1887,  20,  1665).  Thorium  forms  one  oxide,  ThO2  ,  upon  ignition 
of  the  oxalate.  It  is  a  snow-white  powder,  not  easily  soluble  in  acids  if  highly 
ignited  (Cleve,  J.,  1874,  261).  The  hydroxide,  Th(OH)4 ,  is  formed  by  precipita- 
tion of  the  salts  by  the  alkalis.  It  is  a  white,  heavy,  gelatinous  precipitate, 
drying  to  a  hard  glassy  mass.  The  chloride,  ThCl4  ,  and  the  nitrate,  Th(N03)4 , 
*>re  deliquescent.  The  chloride  is  a  white  body  melting  at  a  white  heat  and  then 
subliming  in  beautiful  white  needles  (Kruess  and  Nilson,  I.e.).  The  sulphate 
is  soluble  in  five  parts  of  cold  watiT.  The  carbon-file,  oxalate  and  phosphate  are 
insoluble  in  water;  the  a.ralatc  is  scarcely  soluble  in  dilute  mineral  acids. 
Alkali  hydroxides  or  sulphides  precipitate  thorium  hydroxide,  Th(OH)4 , 
insoluble  in  excess  of  the  reagent.  Tartaric  and  citric  acids  hinder  the  pre- 
cipitation. Alkali  carbonates  precipitate  the  basic  carbonate,  soluble  in  ex- 
cess, if  the  reagent  be  concentrated.  The  solution  in  (NH4)2C03  readily  repre- 
cipitates  upon  warming.  BaC03  precipitates  thorium  salts  completely/  Oxalic 
acid  and  oxalates  form  a  white  precipitate  (distinction  from  Al  and  Gl),  not 
soluble  in  oxalic  acid  or  in  dilute  mineral  acids;  soluble  in  hot  concentrated 
(NH1),C,O4  and  not  reprecipitated  on  cooling  and  diluting  (distinction  from 
Ce  and  La).  A  saturated  solution  of  K2S04  slowly  but  completely  precipitates 
a  solution  of  Th('S04)2  ,  forming  potassium  thorium  sulphate;  insoluble  in  a 
saturated  K,SO4  solution,  sparingly  soluble  in  cold  water,  readily  soluble  in 
hot  water.  HF  precipitates  Th.F4  ,  insoluble  in  excess,  gelatinous,  becoming 
crystalline  on  standing.  Boiling  freshly  precipitated  Th(OH)4  with  KF  in 
presence  of  HF  forms  K2ThF6.4H2O  ,  a  heavy  fine  white  precipitate  almost 
insoluble  in  water.  The  distinguishing  reactions  of  thorium  are  the  precipitation 
with  oxalates  and  with  K2SO4  ,  and  failure  to  form  a  soluble  compound  on 
fusion  with  Na2C03  (distinction  from  Si02  and  Ti02). 

§170.  Titanium.    Ti  =  48.1.    Valence  three  and  four. 

Titanium  is  found  quite  widely  disiribut3d  as  lutiie,  Lrookite,  anatase, 
titanite,  titaniferous  iron,  FeTiOs,  and  in  many  soils  and  clays.  Never  found 
native.  It  is  prepared  by  heating  the  fluoride  or  chloride  with  K  or  Na .  It 
is  a  dark  gray  powder,  which  shows  distinctly  metallic  when  magnified',  melt- 
ing point,  1800°  (Cir.  B.  of  S.,  1915).  Heated  in  the  air  it  burns  with  an  unusu- 
ally brilliant  incandescence',  sifted  into  the  flame  it  burns  with  a  blinding  bril- 
liance. Chlorine  in  the  cold  is  without  action,  when  heated  it  combines  with 
vivid  incandescence.  It  decomposes  water  at  100°.  It  is  soluble  in  acids,  with 
evolution  of  hydrogen,  forming  titanous  salts.  At  a  higher  temperature  it  com- 
bines directly  with  Br  and  I .  It  is  almost  the  only  metal  that  combines 
directly  with  nitrogen  when  heated  in  the  air  (Woehler  and  Deville,  A., 
1857,  103,  230;  Merz,  J.  pr.,  1866,  99,  157).  The  most  common  oxide  of 
titanium  is  the  dioxide,  TiO2 ,  analogous  to  CO2  and  SiO2 .  It  occurs  more 
or  less  pure  in  nature  as  rutile,  brookite  and  anatase;  it  is  formed  by  igni- 
tion of  the  hydrated  titanic  acid  or  of  ammonium  titanate  (Woehler,  J.,  1849, 
268).  Ignition  of  TiO2  in  dry  hydrogen  gives  Ti2O2  ,  an  amorphous  black 
powder,  dissolving  in  H2SO4  to  a  violet-colored  solution  (Ebelmen,  A.  Ch., 
1847,  (3),  20,  392).  TiO  is  formed  when  TiO2  is  ignited  with  Mg:2TiO2  -f- 
Mg  =  TiO  +  MgTiO3  (Winkler,  B.,  1890,  23,  2660).  Other  oxides  have  been 
reported.  Titanic  acid,  TiO2 ,  is  a  white  powder,  melts  somewhat  easier  than 
SiO2 ,  soluble  in  the  alkalis  unless  previously  strongly  ignited.  Mixed  with 
charcoal  and  heated  in  a  current  of  chlorine,  TiCl4  is  formed.  The  bromide 
is  formed  in  a  similar  manner.  TiO2  acts  as  a  base,  forming  a  series  of  stable 
salts;  also  as  an  acid,  forming  titanates.  TiCli  is  a  colorless  liquid,  fuming 
in  the  air;  it  boils  at  136.41°  (Thorpe,  J.  C.,  1880,  37,  329);  it  is  decomposed 
by  water,  forming  titanic  acid,  which  remains  in  solution  in  the  HC1  present. 
Solutions  of  most  of  the  titanic  salts,  when  boiled,  deposit  the  insoluble 
meta-titanic  acid.  HF  dissolves  all  forms^  of  titanic  acid;  if  the  solution 
be  evaporated  in  presence  of  H2SO4  no  TiF.,  is  volatilized  (distinction  from 
SiF4).  When  evaporated  with  HF  alone,  TiF4  is  volatilized.  The  double 
potassium  titanium  fluoride,  K2TiF6 ,  formed  by  fusing  TiO2  with  acid  KF  ,  is 
sparingly  soluble  in  water  (96  parts),  readily  soluble  in  HC1  .  Solutions  of 
titanic  salts  in  water  or  acid  solutions  of  titanic  acid  are  precipitated  by 
alkali  hydroxides,  carbonates  and  sulphides  as  the  hydrated  titanic  acid,  insoluble 
in  excess  of  the  precipitants  and  in  ammonium  salts.  BaCO3  gives  the  same 


206  URANIUM.  §171- 

precipitate.  K4Fe(CN)0  gives  a  reddish-yellow  precipitate;  K,Fe(CN)6  a  yellow 
precipitate.  NaaHPO,  precipitates  the  titanium  ahinjxt  cviuplctcliL  even  in  the 
presence  of  strong-  HC1  .  An  acid  solution  of  TiO,  when  treated  with  Sn  or 
Zn  gives  a  pale  blue  to  violet  coloration  to  the  solution,  due  to  a  partial  reduction 
of  the  titanium  to  the  triad  condition.  These  colored  solutions  are  precipitated 
by  alkali  hydroxides,  carbonates  and  sulphides.  ITS  is  without  action.  The 
solution  reduces  Fe"'  to  Fe"  ,  CM"  to  Cu' ,  and  salts  of  Hg- ,  Ag  and  Au  to  the 
metallic  state;  the  titanium  becoming  again  the  tetrad.  The  reduction  by  Sn 
or  Zn  takes  place  in  presence  of  HF  (distinction  from  columbic  acid).  Titanium 
compounds  fused  in  the  flame  with  microcosmic  salt  give  in  the  reducing  flame 
a  yellow  bead  when  hot,  cooling  to  reddish  and  violet  (reduction  of  the  tita- 
nium). With  FeSOt  in  the  reducing  flame  a  Mood-red  bead  is  obtained. 

Titanium  is  very  readily  detected  in  minerals  as  follows.  0.1  gram  of  the 
finely  powdered  mineral  is  mixed  with  0.2  gram  of  finely  powdered  sodium 
fluoride  and  3  grams  sodium  pyrosulphate  added  without  mixing.  The  crucible 
is  heated  until  copious  sulphuric  acid  fumes  are  evolved.  The  fused  mass  is 
rapidly  cooled  and  heated  with  2-3  c.c.  dilute  sulphuric  acid  and  10  c.c.  water 
added.  The  solution  is  dividqd  into  two  parts  and  a  few  drops  of  hydrogen 
peroxide  added  to  one  part.  A  yellow  color  is  produced  by  the  titanium. 
Chlorides,  bromides  and  iodides  interfere  with  this  very  delicate  reaction  (Weber, 
Z.,  40,  799,  Noyes,  J.  Soc.  Ind.,  10,  485). 

§171.  Uranium,     U  =  238.2.     Valence  four  and  six. 

Specific  gravity,  18.685  (Zimmermann,  A.,  1882,  213,  285).  Melting  point, 
<1850°  (Cir.  B.  of  S.,  1915).  Found  in  various  minerals;  its  chief  ore  is 
pitch-blende,  which  contains  from  40  to  90  per  cent  of  UsOg .  Prepared  by 
fusing  UClt  with  K  or  Na  (Zimmermann,  A.,  1883,  216,  1;  1886,  232,  273). 
It  has  the  color  of  nickel,  hard,  but  softer  than  steel,  malleable,  permanent 
in  the  air  and  water  at  ordinary  temperatures;  when  ignited  burns  with  incan- 
descence to  UgOg  J  unites  directly  with  Cl ,  Br  ,  I  and  S  when  heated;  soluble 
in  HC1 ,  HoSO4  and  slowly  in  HNO3  .  Uranous  oxide,  UO2  ,  formed  by  ignit- 
ing the  higher  oxides  in  carbon  or  hydrogen,  is  a  brown  powder,  soon  turning 
yellow  by  absorption  of  oxygen  from  the  air.  Uranous  hydroxide  is  formed 
by  precipitating  uranous  salts  with  alkalis.  Uranic  oxide,  UOs  ,  is  formed 
by  heating  uranic  nitrate  cautiously  to  25°,  and  upon  ignition  in  the  air  both 
this  and  other  uranium  oxides,  hydroxides  and  uranium  oxysalts  with  volatile 
acids  are  converted  into  U3O8  =  UO22UO3 .  Uranium  acts  as  a  base  in  two 
classes  of  salts,  uranous  and  uranyl  salts.  Uranous  salts  are  green  and  give 
green  solutions,  from  wyhich  alkalis  precipitate  uranous  hydroxide,  insoluble  in 
excess  of  the  alkali;  alkali  carbonates  precipitate  U(OH)4,  soluble  in 
(NH4)2CO3;  with  BaCO3  the  precipitation  is  complete  even  in  the  cold.  H2S  is 
without  action;  (NH4)J3  gives  a  dark-brown  precipitate;  K4Fe(CN)6  gives  a 
reddish-brown  precipitate.  In  their  action  toward  oxidizing  and  reducing 
agents  uranous  and  uranyl  (uranic)  salts  resemble  closely  ferrous  and  ferric 
salts;  uranous  salts  are  even  more  easily  oxidized  than  ferrous  salts,  e.  g.,  by 
exposure  to  the  air,  by  HNO3 ,  Cl ,  HC1O3 ,  Br  ,  KMnO4 ,  etc.  Gold,  silver  and 
platinum  salts  are  reduced  to  the  free  metal.  The  hexad  uranium  (U^1)  acts 
as  a  base,  but  usually  forms  basic  salts,  never  normal:  we  have  TJ02(NO:J)2 , 
not  TT(NO3)0;  UOLS04 ,  not  TT($O4)3  .  These  basic  salts  wTere  formerly  called 
uranic  salts,  but  at  present  (ir02)"  is  regarded  as  a  basic  radical  and  called 
tirunjil,  and  its  salts  are  called  uranyl  salts,  e.g.,  UO2C12  uranyl  chloride, 
(TTO2)S(PO4)2  uranyl  orthophosphate.  Solutions  of  uranyl  salts  are  yellow; 
KOH  and  NaOH  give  a  yellow  precipitate,  uranates,  K2TJ207  and  Na.ILO,  , 
insoluble  in  excess.  Alkali  carbonates  give  a  yellow  precipitate,  soluble  in 
excess;  BaC03  and  CaC03  give  TJ03 .  H2S  does  not  precipitate  the  uranium, 
but  slowly  reduces  uranyl  salts  to  uranous  salts  (Formanek,  A.,  1890,  257,  115)1 
(NH4)..S  gives  a  dark-brown  precipitate.  K4Fe(CN)0  gives  a  reddish-brown 
precipitate.  Used  in  the  analysis  and  separation  of  uranium  compounds 
(Fresenius  and  Hintz,  Z.  angeic.,  1895,  502).  Sodium  phosphate  gives  a  yellow 
precipitate.  The  hexad  uranium  acts  as  an  acid  toward  some  stronger  bases. 


§171  rt,  8.  VANADIUM-  207 


Thus  we  have  K2TT2O7  and  Na^U^Oy  ,  formed  by  precipitating  uranyl  salts  with  KOH 
and  NaOH  ;  compare  the  similar  salts  of  the  hcxad  chromium,  K2Cr2O7  and 
Na2Cr2O7  .  Other  oxides  of  uranium  are  described,  but  are  doubtless  combinations 
of  UO2  and  UO3  .  Zn  ,  Cd  ,  Sn  ,  Pb  ,  Co  ,  Cu  ,  Fe  ,  and  ferrous  salts  reduce  uranyl 
salts  to  uranous  salts.  Solutions  of  Sn,  Ft,  Au,  Cu,  Hg  and  Ag  are  reduced  to 
the  metal  by  metallic  uranium  (Zimmcrmann,  I.e.).  For  method  of  recovery  of 
waste  uranium  compounds,  see  Laube  (Z.  angew.,  1889,  575). 

§171a.     Vanadium.     V  =  51.0.     Valence  two  to  five. 

1.  Properties.—  -Specie  gravity,  5.8  (Moissan,  C.  r.,  122);    melting  point   1720° 
(Cir.  B.  S.,  35,  1915).     A  grayish  non-magnetic  powder;   slowly  oxidized  in  the  air, 
rapidly   on  ignition  with  formation  of  V2O5  .     It  forms  with  chlorine  the  dark 
brown  tetrachloride. 

2.  Occurrence.  —  It  is  often  found  in  iron  and  copper  ores  and  in  some  clays  and 
rare  minerals,  e.g.,  vanadinite,    (Pb5Cl(VO4)3)  ;    volborthite,  (Cu,Ca,Ba)3(OH)3VO4 
+6H2O)  ;      mottramite      (a    hydrous     vanadate    of    lead,      copper,     and     other 
divalent  elements,  of  uncertain  formula,   (R3(VO4)2.3R(OH)2))  . 

3.  Preparation.  —  The  vanadium  ores  are  treated  chiefly  for  the  preparation  of 
ammonium  vanadate  and  vanadic  acid.     The  ores  are  fused  with  KN03  ,  form- 
ing- potassium  vanadate.     This  is  precipitated  with  Pb  or  Ba  salts  and   then 
decomposed  with  H.SO4  .     The  vanadic  acid  is  neutralized  with  NH,OH  and 
precipitated  with  NH4C1  ,  in  which  it  is  insoluble.     This  upon  ignition  gives 
V,O0  pure  (\Yohler,  A.,  1851,  78,  125).     The  metal  is  prepared  from  the  dichlo- 
ride,  VCL  ,  by  long-continued  ignition  in  a  current  of  hydrogen. 

4.  Oxides.  —  Vanadium  forms  four  oxides:  VO  ,  gray;  V203  ,  black;  V02  ,  dark 
blue;  and  V,O,,  ,  dark  red  to  orange  red. 

5.  Solubilities.  —  Vanadium  is  not  attacked  by  dilute  HC1  or  H2SO4;  concen- 
trated H,SO4  gives  a  greenish-yellow  solution;  HNO3  a  blue  solution.     VO  dis- 
solves in  acids  to  a  blue  solution  with  evolution  of  hydrogen.     V2O3  dissolves 
in  dilute  HC1  to  a  dark  greenish-black  solution.     Chlorine  forms  with  V2O3  , 
VOC13  and  V2Or,  .     VO.,  dissolves  in  acids  to  a  blue  solution,  from  which  solu- 
tions Na.CO,   gives  a  precipitate  of  V202(OH)4   +  5H2O  ,  grayish-white  mass, 
losing   4H2O   at    100°    and   turning  black,   soluble   in   acids    and   alkalis.     V2O5 
exists    in    several   modifications   with    different   solubilities   in   water,    the   red 
modification  being-  soluble  in  125  parts  of  water  at  20°   (Ditte,  C.  r.,  1880,  101, 
698).     Vanadic  acid  forms  three  series  of  salts,  ortho,  meta  and  pyro,  analogous 
to  the  phosphates.     Most  salts  are  the  metavanadates.     The  ortho  compounds 
are  quite  unstable,  readily  changed  to  the  meta  and  pyro  compounds.     Alkali 
vanaclates  are  soluble  in  water,  the  ammonium  vanadate  least  soluble  and  not 
at  all  in  NH4C1  . 

6.  Reactions.  —  Solutions    of   vanadic   acid   produce   brown   precipitates    with 
alkalis,  soluble  in  excess  to  a  yellowish-brown  color.     Potassium  ferrocyanide 
gives  a  green  precipitate,  insoluble  in  acids.     Tannic  acid  gives  a  blue-black 
solution,  which  is  said  to  make  a  desirable  ink.     Ammonium  sulphide  precipi- 
tates V,S5  ,  brown,  soluble  with  some  difficulty  in  excess  of  the  reagent  to  a 
reddish-brown   thio   salt.     From   this   solution    acids    reprecipitate    the   brown 
vanadic  sulphide,  V2S5  . 

If  to  a  solution  of  a  vanadate,  neutral  or  alkaline,  solid  NH4C1  be  added,  the 
vanadium  is  completely  precipitated  as  NH4V03  ,  ammonium  metavanadate, 
crystalline,  colorless,  insoluble  in  NH4C1  solution;  upon  ignition  in  air  or  oxy- 
gen, pure  vanadic  oxide,  V20B  ,  is  obtained. 

7.  Ignition.  —  Borax  gives  with  vanadium  compounds  in  the   outer  flame  a 
colorless  bead,  yellow  if  much  vanadium  be  present;  in  the  inner  flame  a  green 
bead,  or  browrn  when  vanadium  is  present  in  large  quantities  and  hot,  becoming 
green    upon    cooling.     All    the    lower   oxides    of    vanadium    ignited    in    air    or 
oxygen  give  V2O5  . 

8.  Detection.—  Vanadium  will  almost  always  be  found  as  a  vanadate  (2)  and 
is  detected  by  the  reactions  used  in  its  purification  (3)  ;  also  by  the  reactions 
with  reducing  agents,  forming-  the  colored  lower  oxidized  compounds  (10). 


208  YTTERBIUM—YTTRIUM.  §171  rt,  10. 

9.  Estimation. — (1}     It   is  precipitated  as  basic  lead  vanadate  and  dried   at 
100°.      (2}     It  is    precipitated    as    ammonium    vanadate,     NH4VO3  ,     in    strong 
NH4C1  solution,  ignited  to  the  oxide  V-Oa  ,  and  weighed. 

10.  Oxidation. — Zn  ,  in  solutions  of  vanadates  with  dilute  H2SOi  ,  reduces  the 
vanadium  to  the  tetrad,  a  green  to  blue  solution,   then  greenish-blue  to  green, 
the  triad,  and  finally  to  lavender  blue,  the  dyad.     H2S  reduces  vanadates  to  the 
tetrad  with  separation  of  sulphur.     Oxalic  acid  and  sulphurous  acid  also  reduce 
vanadates  to  the  tetrad,  the  solution  becoming  blue. 

§172.  Ytterbium.     Yb  =  173.5.     Valence  three. 

Obtained  as  an  earth  by  Marignac  (C.  r..  1878,  87,  578)  from  a  gadolinite 
earth;  by  Delafontaine  (C.  r.,  1878,  87,  933)  from  sipylite  found  at  Amherst,  Va. 
Nilson  (/?.,  1879,  12,  550;  1880,  13,  1433)  describes  its  preparation  from  euxenite 
and  its  separation  from  Sc  .  It  has  the  lowest  bacisity  of  the  yttrium  earths. 
The  double  potassium  ytterbium  sulphate  is  easily  soluble  in  water  and  in 
potassium  sulphate.  The  oxalate  forms  a  white  crystalline  precipitate,  in- 
soluble in  water  and  in  dilute  acids.  The  salts  are  colorless  and  give  no 
absorption  spectrum.  For  the  spark  spectrum  see  Welsbach  (J/.,  1884,  5,  1). 
The  oxide,  Yb,,03  ,  is  a  white  powder,  slowly  soluble  in  cold  acids,  readily  upon 
warming.  The  Jiydro.riilc  forms  a  gelatinous  precipitate,  insoluble  in  i^H4OH 
but  soluble  in  KOH  .  It  absorbs  CO,  from  the  air.  The  nitrate  melts  in  it« 
water  of  crystallization  and  is  very  soluble  in  water. 

§173.  Yttrium.     Y  =  88.7.    Valence  three. 

Yttrium  is  one  of  the  numerous  rare  metals  found  in  the  gadolinite  mineral 
at  Ytterby,  near  Stockholm,  Sweden:  also  found  in  Colorado  (Hidden  and 
Mackintosh,  Am-.  £'.,  1889,  38,  474).  The  metal  has  been  prepared  by  electro- 
lysis of  the  chloride;  also  by  heating  the  oxide,  Y203  ,  with  Mg-  (Winkler,  B., 
1890,  23,  787).  Melting  point,  1490°  (Cir.  B.  of  S.~  1915).  The  study  of  these 
rare  earths  is  by  no  means  complete.  It  is  also  claimed  that  they  have  not 
yet  been  obtained  pure,  but  that  the  so-called  pure  oxides  really  consist  of  a 
mixture  of  oxides  of  from  five  to  twenty  elements  (Crookes,  C.  N.,  1887,  55, 
107,  119  and  131).  The  most  of  these  rare  earths  do  not  give  an  absorption 
spectrum,  but  give  characteristic  spark  spectra;  and  it  is  largely  by  this  means 
that  the  supposedly  pure  oxides  have  been  shown  to  be  mixtures  of  the  oxides 
of  several  closely  related  elements  (Welsbach,  M.,  1883,  4,  641;  Dennis  and 
Chamot,  J.  Am.  Soc.,  1897,  19,  799).  Yttrium  salts  are  precipitated  by  the 
ai  -alis  and  by  the  alkali  sulphides  as  the  hydroxide,  Y(OH)?  ,  a  white  bulky  pre- 
cipitate, insoluble  in  the  excess  of  the  reagents  (distinction  from  Gl).  The 
oxide  and  hydroxide  are  readily  soluble  in  acids;  boiling  with  NH,C1  causes 
solution  of  the  hydroxide  as  the  chloride.  The  alkali  carbonates  precipitate 
the  carbonate  Y2(CO3)3 ,  soluble  in  a  large  excess  of  the  reagents.  If  the  solu- 
tion in  ammonium  carbonate  be  boiled,  the  hydroxide  is  precipitated.  Soluble 
oxalates  precipitate  yttrium  salts  as  the  white  oxalate  (distinction  from  Al 
and  Gl);  soluble  with  some  difficulty  in  HC1  .  The  double  sulphate  with 
potassium  is  soluble  in  water  and  in  potassium  sulphate  (distinction  from 
thorium,  zirconium  and  the  cerite  metals).  BaCO3  forms  no  precipitate  in 
the  cold  (distinction  from  Al  ,  Fe'"  ,  Cr'"  ,  Th  ,  Ce  ,  La  ,  Nd  and  Pr).  Hydro- 
fluoric acid  precipitates  the  gelatinous  fluoride,  YF3  ,  insoluble  in  water  and 
in  HF  .  The  precipitation  of  yttrium  salts  is  not  hindered  by  the  presence 
of  tartaric  acid  (distinction  from  Al  ,  Gl  ,  Th  and  Zr).  The  analysis  of  yttrium 
usually  consists  in  its  detection  and  separation  in  gadolinite  (silicate  of  Y  , 
Gl  ,  Fe  ,  Mn  ,  Ce  and  La).  Fuse  with  alkali  carbonate,  decompose  with  KC1  , 
and  filter  from  the  SiO2 .  Neutralize  the  filtrate  and  precipitate  the  Y  ,  La 
and  Ce  as  oxalates  with  (NH4)2C2O1  .  Ignite  the  precipitate  and  dissolve  in 
HC1  .  Precipitate  the  La  and  Ce  as  the  double  potassium  sulphates,  and  from 
the  filtrate  precipitate  the  yttrium  as  the  hydroxide  with  NH,OH .  Ignite 
and  weigh  as  the  oxide.  In  order  to  effect  complete  separations  the  operations 
should  be  repeated  several  times. 


§175.  TffE  CALCIUM  GROUP.  £09 

§174.  Zirconium.     Zr  =  90. G.     Valence  four. 

Zirconium  is  a  rare  metal  found  in  various  minerals,  chiefly  in  zircon, 
a  silicate;  never  found  native.  The  metal  was  first  prepared  by  Berzelius 
in  1824  by  fusion  of  the  potassium  zirconium  fluoride  with  potassium  (Pogg., 
1825,  4,  117).  Also  prepared  by  electrolysis  of  the  chloride  (Becquerel,  A. 
CA,  1831,  48,  337).  Melting  point,  1700°  ?  (Cir.  B.  of  S.,  1915).  The  metal 
exists  in  three  modifications:  crystalline,  graphitoidal  and  amorphous.  The 
amorphous  zirconium  is  a  velvet-black  powder,  burning  when  heated  in  the 
air.  Acids  attack  it  slowly  even  when  hot,  except  HF  ,  which  dissolves  it  in 
the  cold.  It  forms  but  one  oxide,  ZrO2 ,  analogous  to  SiO2  and  TiO> .  ZrO2 
is  prepared  from  the  mineral  zircon  by  fusion  with  a  fixed  alkali.  Digestion 
in  water  removes  the  most  of  the  silicate,  leaving  the  alkali  zirconate  as  a 
sandy  powder.  Digestion  with  HC1  precipitates  the  last  of  the  SiOa  and 
dissolves  the  zirconate.  The  solution  is  neutralized,  strongly  diluted  and 
boiled,  whereupon  the  zirconium  precipitates  as  the  basic  chloride  free  from 
iron.  Or  the  zirconium  may  be  precipitated  by  a  saturated  solution  of  K2SO4  , 
and  after  resolution  in  acids  precipitated  by  NH4OH  and  ignited  to  ZrO_> 
(Berlin,  J.  pr.,  1853,  58,  145;  Roerdam,  C.  C.,  1889,  533).  ZrO2  is  a  white 
infusible  powder,  giving  out  an  intense  white  light  when  heated;  it  shows  no 
lines  in  the  spectrum.  It  is  much  used  with  other  rare  earths,  La.oO;(  ,  Y2O3  , 
etc.,  to  form  the  mantles  used  in  the  Welsbach  gas-burners  (Drossbach,  C.  C., 
1891,  772;  Welsbach,  </.,  1887,  2670;  C.  N.,  1887,  55,  192).  The  oxide  (or  hydroxide 
precipitated  hot)  dissolves  with  difficulty  in  acids  to  form  salts.  The  hydroxide, 
ZrO(OH)2  ,  precipitated  in  the  cold  dissolves  readily  in  acids.  As  an  acid, 
zirconium  hydroxide,  ZrO(OH)2  =  H2ZrO3  ,  forms  zirconates,  decomposed  by 
acids.  As  a  base  it  forms  zirconium  salts  with  acids.  The  sulphate  is  easily 
soluble  in  water,  crystallizing  from  solution  with  4H.O  .  The  phosphate  is 
insoluble  in  water,  formed  by  precipitation  of  zirconium  salts  by  Na2HP04  or 
H3PO4  .  The  silicate,  ZrO2.SiO2  ,  is  found  in  nature  as  the  mineral  zircon, 
usually  containing  traces  of  iron.  Zirconium  chloride  is  formed  when  a  current 
of  chlorine  is  passed  over  heated  ZrO2  .  mixed  with  charcoal.  It  is  a  white 
solid,  may  be  sublimed,  is  soluble  in  water.  Solutions  of  zirconium  salts  are 
precipitated  as  the  hydroxide,  ZrO(OH)2  ,  by  alkali  hydroxides  and  sulphides, 
a  white  flocculent  precipitate,  insoluble  in  excess  of  the  reagents,  insoluble  in 
NH4C1  solution  (difference  from  Gl).  Tartaric  acid  prevents  the  precipitation. 
Alkali  carbonates  precipitate  basic  zirconium  carbonate,  white,  soluble  in 
excess  of  KHCO3  or  (NH4)2C03;  boiling  precipitates  a  gelatinous  hydroxide 
from  the  latter  solution.  BaCO3  does  not  precipitate  zirconium  salts  com- 
pletely, even  on  boiling.  The  precipitates  of  the  hydroxide  and  carbonate  are 
soluble  in  acids.  Oxalic  acid  and  oxalates  precipitate  zirconium  oxalate,  solu- 
ble in  excess  of  oxalic  acid  on  warming,  and  soluble  in  the  cold  in  (NH4)2C204 
(difference  from  thorium);  soluble  in  HC1  .  A  saturated  solution  of  K2S04 
precipitates  the  double  potassium  zirconium  sulphate,  white,  insoluble  in  excess 
of  the  reagent  if  precipitated  cold,  soluble  in  excess  of  HC1;  if  precipitated 
hot,  almost  absolutely  insoluble  in  water  or  HC1  (distinction  from  Th  and  Ce). 
Zirconium  salts  are  precipitated  on  warming  with  Na,S.,O3  (separation  from 
Y,  Nd  and  Pr).  Solution  of  H,O,  completely  precipitates  zirconium  salts. 
Tumcric  paper  moistened  with  a  solution  of  zirconiiim  salt  and  HC1  is  colored 
orange  upon  drying  (boric  acid  gives  the  same  reaction)  (Brush,  J.  pr.,  1854, 
62,  7).  HF  does  not  precipitate  zirconium  solutions,  as  zirconium  fluoride, 
ZrF4  ,  is  soluble  in  water  and  in  HF  (distinction  from  Th  and  Y). 


THE  CALCIUM  GROUP  (FIFTH  GROUP). — (THE  ALKALINE  EARTH  METALS.) 

Barium.    Ba  =  137.37.     Calcium.     Ca  =  40.07. 

Strontium.     Sr  =  87.03.     Magnesium.     Mg  =  24.32. 

§175.    Like  the  alkali  metals,  Ba ,  Sr ,  and  Ca  oxidize  rapidly  in  the  ai- 

at  ordinary  temperatures — forming  alkaline  earths — and  dcnniijtone  wuft  >\ 

forming  hydroxides  with  evolution  of  heat.     Mg  oxidizes  rapidly  in  the  air 


210  THE  CALCIUM  GROUP.  §176. 

when  ignited,  decomposes  water  at  100°,  and  its  oxide — in  physical  proper- 
ties farther  removed  from  Ba  ,  Sr ,  and  Ca  than  these  oxides  are  from  e;icli 
other — slowly  unites  with  water  without  sensible  production  of  heat.  As 
compounds,  these  metals  are  not  easily  oxidized  beyond  their  quantivalence 
as  dyads,  and  they  require  very  strong  reducing  agents  to  restore  them 
to  the  elemental  state. 

§176.  In  basic  power,  Ba  is  the  strongest  of  the  four,  Sr  somewhat 
stronger  than  Ca,  and  Mg  much  weaker  than  the  other  three.  It  will  be 
observed  that  the  solubility  of  their  hydroxides  varies  in  the  same  decreas- 
ing gradation,  which  is  also  that  of  their  atomic  weights;  while  the 
solubility  of  their  sulphates  varies  in  a  reverse  order,  as  follows:  (§7) : 

§177.  The  hydroxide  of  Ba  dissolves  in  about  30  parts  of  water;  that  of 
Sr,  in  100  parts;  of  Ca,  in  800  parts;  and  of  Mg,  in  100,000  parts.  The 
sulphate  of  Ba  is  not  appreciably  soluble  in  water  (429,700  parts  at  18.4°; 
Hollemann,  Z.  phys.  Ch.,  1893,  12,  131);  that  of  Sr  dissolves  in  10,000 
parts;  of  Ca ,  in  500  parts;  of  Mg ,  in  3  parts.  To  the  extent  in  which  they 
dissolve  in  water,  alkaline  earths  render  their  solutions  caustic  to  the 
taste  and  touch,  and  alkaline  to  test-papers  and  phenolphthalein. 

§178.  The  carbonates  of  the  alkaline  earths  are  not  entirely  insoluble 
in  pure  water:  BaC03  is  soluble  in  45,566  parts  at  24.2°  (Hollemann, 
Zeit.  phys.  Ch.,  1893,  12,  125);  SrC03  in  90,909  parts  at  18°  (Kohlrausch 
and  Rose,  Zeit.  phys.  Ch.,  1893,  12,  241);  CaC03  in  80,040  parts  at  23.8° 
(Hollemann,  /.  c.);  MgC03  in  9,434  parts  (Chevalet,  Z.,  1869,  8,  91).  The 
presence  of  NH4OH  and  (NH4)2C03  lessens  the  solubility  of  the  carbonates 
of  Ba ,  Sr ,  and  Ca ,  while  their  solubility  is  increased  by  the  presence  of 
NH4C1 .  MgC03  is  soluble  in  ammonium  carbonate  and  in  ammonium 
chloride,  so  much  so  that  in  presence  of  an  abundance  of  the  latter  it  is 
not  at  all  precipitated  by  the  former,  i.  e.  (NH4).,C03  does  not  precipitate  a 
solution  of  MgCl2  as  the  NH4C1  formed  holds  the  Mg  in  solution. 

§179.  These  metals  may  be  all  precipitated  as  phosphates  in  presence 
of  ammonium  salts,  but  their  further  separation  for  identification  or  esti- 
mation would  be  attended  with  difficulty  (§145  and  //.). 

§180.  The  oxalates  of  Ba ,  Sr,  and  Mg  are  sparingly  soluble  in  water, 
calcium  oxalate  insoluble.  Barium  chromate  is  insoluble  in  water  (§§27 
and  186,  5^),  strontium  chromate  sparingly  soluble,  and  calcium  and  mag- 
nesium chromates  freely  soluble. 

§181.  In  qualitative  analysis,  the  group-separation  of  the  fifth-group 
metals  is  effected,  after  removal  of  the  first  four  groups  of  bases,  by 
precipitation  with  carbonate  in  presence  of  ammonium  chloride,  after 
which  magnesium  is  precipitated  from  the  filtrate,  as  phosphate. 

§182.  The  hydroxides  of  Ba,  Sr,  and  Ca,  in  their  saturated  solutions, 
necessarily  dilute,  precipitate  solutions  of  salts  of  the  metals  of  the  first 


§186,  4.  BARIUM.  211 

four  groups  and  of  Mg  ,  as  hydroxides.  In  turn,  the  fixed  alkalis  precipi- 
tate, from  solutions  of  Ba ,  Sr ,  Ca ,  and  Mg ,  so  much  of  the  hydroxides 
of  these  metals  as  does  not  dissolve  in  the  water  present*;  hut  ammonium 
hydroxide  precipitates  only  Mg ,  and  this  but  in  part,  owing  to  the  solubility 
of  Mg(OH)2  in  ammonium  salts. 

§183.  Solutions  containing  Ba  ,  Sr ,  Ca  ,  and  Mg ,  and  phosphoric,  oxalic, 
boric,  or  arsenic  acid,  necessarily  have  an  acid  reaction,  because  these 
phosphates,  oxalates,  etc.,  are  soluble  only  in  acids;  such  solutions  are 
precipitated  by  ammonium  hydroxide  or  by  any  agent  which  neutralizes  the 
solution,  and,  consequently,  we  have  precipitates  of  this  kind  in  the  third 
group  (§145  and  //.)  : 

CaClo  +  H,PO4  +  2NH4OH  =  CaHPO4  +  2NH4C1  +  2H2O 
CaH4(PO4)2  +  2NH4OH  =  CaHPO4  +  (NH4)2HPO4  +  2H2O  . 

If  excess  of  the  ammonium  hydroxide  be  added  the  precipitate  is  Ca3  (P04)9 . 
Barium  and  strontium  react  like  calcium.  In  the  case  of  a  magnesium 
bait  the  precipitate  is  MgNH4P04 . 

§184.  The  carbonates  of  the  alkaline  earth  metals  are  dissociated  by 
heat,  leaving  metallic  oxides  and  carbonic  anhydride.  This  occurs  only  at 
a  high  temperature  in  the  case  of  Ba . 

§185.  Compounds  of  Ba  ,  Sr  ,  and  Ca  (preferably  with  HC1)  impart  char- 
acteristic colors  to  the  non-luminous  flame,  and  readily  present  well-defined 
spectra. 

§186.    Barium.     Ba=l37.37.     Valence  two. 

1.  Properties.—  Specific  gravity,   3.75   (Kern,   C.   N.,    1875,   31,  243);    melting 
point,  850°   (Cir.  B.  of  £.,   1915).     It  is  a  white  metal,  stable  in  dry  air,  but 
readily  oxidized  in  moist  air  or  in  water  at  ordinary  temperature,  hydrogen  being 
evolved  and  barium  hydroxide  formed.     It  is  malleable  and  ductile  (Kern,  I.  c.). 

2.  Occurrence. — Barium  can  never  occur  in  nature  as  the  metal  or  oxide,  or 
hydroxide  near  the  earth's  surface,   as  the  metal  oxidizes  so  readily,   and  the 
oxide   and   hydroxide   are   so   basic,   absorbing  acids  readily  from   the  air.     Its 
most  common  forms  of  occurrence  are  barite,  BaSO4 ,  and  witherite,  BaCOj . 

3.  Preparation. — (1)  By  electrolysis  of  the  chloride  fused  or  moistened  with 
strong-  HC1 .     (2)  By  electrolysis  of  the  carbonate,  sulphate,  etc.,  mixed  with 
Hg  and  HgO  ,  and  then  distilling-  the  amalgam.     (3)  By  heating-  the  oxide  or 
various  salts  with  sodium  or  potassium  and  extracting1  the  metal  formed  with 
mercury,  then  separating-  by  distillation  of  the  amalgam. 

4.  Oxides  and  Hydroxides. — The  oxide,  BaO  ,  is  formed  by  the  action  of  heat 
upon  the  hydroxide,  carbonate,  nitrate,  oxalate,  and  all  its  organic  salts.     The 
corresponding  hydroxide,  Ba(OH)2  ,  is  made  by  treating  the  oxide  with  water. 
The  peroxide,  Ba<X  ,  is  made  by  heating  the  oxide  almost  to  redness  in  oxygen, 
or  air  which  has  been  freed  from  carbon  dioxide;  by  heating  the  oxide  with 
potassium  chlorate  (Liebig,  Fogg.,  1832,  26,  172)  or  cupric  oxide  (Wanklyn,  7?., 
1874,  7,  1029).     It  is  used  as  a  source  of  oxygen,  which  it  gives  off  at  a  white 
heat,  BaO  remaining;  also  in  the  manufacture  of  hydrogen  peroxide,   H...O2  , 
which  is  formed  by  treating  it  with  dilute  acids:  BaO2   +  2HC1  =  BaCl2   + 
H202. 

*The  presence  of  an  excess  of  fixed  alkali  renders  these  hydroxides  much  less  soluble,  the 
high  concentration  of  the  hydroxyl  ions,  one  of  the  factors  of  the  solubility  product,  diminish- 
ing the  other  factor.  (§45). 


XM-:  BARIUM.  §186, 5a. 

,"».  Solubilities. — rt. — Metal. —  Motallio  bnvimn  is  ivadilv  soluble  in  ncids  with 
evolution  of  hyilro^vn.  f>.  ().n'(/r,v  ami  Iun1ro.ri<1c>i.-  Harinm  oxide  is  aeted  upon 
b\  \\ater  \\ith  evolution  of  heat  .iud  formation  of  the  hydroxide,  \\hieh  is 
soluble  in  about  .'(>  parts  of  cold  \\ater  and  in  its  oxvn  weight  of  hot  water 
(  KVxiMislhfil  and  Rnehlinann.  •  /.,  1S70.  :ni).  Hariuiu  peroxide,  BaO  .  is  very 
sparingly  soluble  in  \\ater  (Sehone.  .1.,  is;?.  192.  857);  soluble  in  aeids  with 
formation  of  H,O_,  . 

r. — Salts. — Most  of  the  soluble  sails  of  barium  are  permanent  :  the 
acetate  is  o!llorcseont.  rriio  elilorido,  bromide.  l>romnto.  iodide.  sul]>hide. 
t\MT(H'vanide,  nitrato.  liypo])bosphito,  oblorato.  neetnie,  niul  pbenvlsul- 
]>bate,  ;in«  freely  soluble  in  water:  the  carbonate.  *u!j>liatc.  sulphite, 
chronuilr,*  |u1ios]>hite.  jihosphate,  oxalate.  iodate,  and  sitico-fluoride,  are 
insoluble  in  water.  The  sulphate  is  perceptibly  soluble  in  strong  HC1  . 
The  chloride  is  almost  insoluble  in  strong  hydrochloric  acid  (separation 
from  Ca  and  Mg)  (Mar.  Am.  $„  1890,  143,  .VJ1V,  likewise  the  nitrate  in 
strong  hydrochloric  and  nitric  acids.  The  chloride  and  nitrate  are  insolu- 
ble in  alcohol. 

().  Reactions,  a. — The  fixed  alkali  hydroxides  precipitate  only  con- 
centrated solutions  of  barium  salts  ("iM.  No  precipitate  is  formed  with 
ammonium  hydroxide  (§45).  The  alkali  carbonates  precipitate  barium 
carbonate.  BaCO,  ,  white.  The  precipitation  is  promoted  by  heat  and 
by  ammonium  hydroxide,  but  is  made  slightly  incomplete  by  the  presence 
of  ammonium  salts  (Yogel.  J.  /'/•..  lSTn>.  7,  1,V>). 

Barium  Carbonate — BaCO, — is  a  valuable  reagent  for  special  -pur 
chietly  for  separation  of  third  and  fourth  group  metals.  It  is  used  in  the 
form  of  the  moist  precipitate,  which  must  be  thoroughly  washed.  It  is 
best  precipitated  from  boiling  solutions  of  barium  chloride  and  sodium  or 
ammonium  carbonate,  washed  once  or  twice  by  decantation.  then  by  tilt  ra- 
tion, till  the  washings  no  longer  precipitate  solution  of  silver  nitrate. 
Mixed  with  water  to  consistence  of  cream,  it  may  be  preserved  for  some 
time  in  stoppered  bottles,  being  shaken  whenever  required  for  use.  When 
dissolved  in  hydrochloric  arid,  and  fully  precipitated  by  sulphuric  acid, 
the  tilt  rate  must  yield  no  fixed  residue.  This  reagent  removes  sulphuric 
acid  (radical)  from  all  sulphates  in  solution  to  which  it  is  added  (c):  Na.,S04 
-f  BaCO.,  ==  BaS04  -f  Na,CO,  .  When  salts  of  non-alkali  metals  are  so 
decomposed,  of  course,  they  are  left  insoluble,  as  carbonates  or  hydroxide-, 
nothing  remaining  in  solution: 

FeSO,  +  BaCOs  =  BaS04  -f  FeCOs 

Fe,(S04)s  -f  ;.BaCO:i  +  :;H,0  —  ;;BaSO4  -f  2Fe(OHV,  +  :;CO, 

The  chlorides  of  the  third  group,  except  Fe"  ,  are  decomposed  by  barium 
carbonate;  while  the  metals  of  the  fourth  group  (/inc.  manganese,  cobalt, 
nickel),  are  not  precipitated  from  their  chlorides  by  this  reagent.  Tartaric 

*Kohlrausob  and  Hose.  Z.  j>/ij/s.  ('/«.,  1S'.V>,  12,  241 :  Schwoit/er,  '/..,  18W,  29,  414. 


£186,  7.  BARIUM.  213 

<cid,  citric  acid,  sugar,  and   other  organic  substances,  hinder  or  prevent' 
'.he  decomposition   by   barium  carbonate. 


1). — Ammonium  oxalate  precipitates  barium  oxalatc,  BaCv04  ,  from  solutions 
of  barium  sails,  sparingly  soluble  in  water,  more  soluble  in  presence  of  am- 
monium chlorldej  soluble  in  oxalic  and  acetic  acids  (Souchay  and  Lciissen,  1., 
1  «:,<;,  99,  36). 

('. — Solutions  of  barium  salts  are  precipitated  by  the  addition  of  concentrated 
nitric  acid  (fir),  rf.— Soluble  phosphates,  full  metallic,  or  two-thirds  metallic, 
as  Na._,HPO,  ,  precipitate  barium  phosphate,  while,  consisting  of  BaHPO, 
when  the  reagent  is  two-thirds  metallic,  and  Ba^FO.,).,  when  1  he  reagent  is 
full  metallic.  Soluble  phosphites  precipitate  barium  salts,  hypophosphites  do 
not.  f.  Mariiim  sulphide  is  not  formed  in  the  wet  way,  hence  hydrosulphuric 
acid  and  soluble  sulphides  are  without  action  upon  barium  salts.  Soluble 
sulphites  precipitate  solutions  of  barium  salts  as  barium  sulphite,  BaSO:,  ,  in- 
soluble in  water  but  soluble  in  hydrochloric  acid  (distinction  from  sulphates). 

Sulphuric  acid,  H2S04 ,  and  all  soluble  sulphates,  precipitate  "barium 
sulphate  (BaSOJ,  white,  very  soluble  in  hot  concentrated  sulphuric 
acid.  Immediate  precipitation  by  the  (dilute  §188,  5c)  saturated  solution 
of  calcium  sulphate  distinguishes  Ba  from  Sr  (and  of  course  from  Ca) ;  but 
precipitation  by  the  (very  dilute  §187,  5r)  solution  of  strontium  sulphate 
is  a  more  certain  test  between  Ba  and  Sr .  BaS04  is  not  transposed  by 
solutions  of  alkali  carbonates  (distinction  from  Sr  and  Ca,  §188,  6a  foot- 
'note). 

/". — Solutions  of  iodates,  as  NaI03  ,  precipitate,  from  barium  solutions  not 
very  dilute,  barium  iodate,  Ba(IO3)2  ,  white,  soluble  in  600  parts  of  hot  or 
17  Hi  parts  of  cold  water  (distinction  from  the  other  alkaline  earth  metals). 
</.  Xeut  ral  or  ammoniacal  solutions  of  arsenous  acid  do  not  precipitate  barium 
salts  (distinction  from  calcium).  Soluble  arsenates  precipitate  solutions  of 
barium  salts,  soluble  in  acids,  including  arsenic  acid. 

JL — Soluble  chromates,  as  K,Cr04 ,  precipitate  solutions  of  barium  salts 
as  barium  chromate,  BaCr04 ,  yellow;  almost  insoluble  in  water  (separa- 
tion from  calcium  and  from  strontium  except  in  concentrated  solutions), 
sparingly  soluble  in  acetic  acid,  moderately  soluble  in  chromic  acid  and 
readily  soluble  in  hydrochloric  and  nitric  acids.  Bichromates,  as  K2Cr207 , 
precipitate  solutions  of  barium  salts  (better  from  the  acetate)  as  the 
normal  chromate  (very  accurate  separation  from  strontium  and  calcium) 
(Grittner,  Z.  angew.,  1892,  73). 

i. — Fluosilicic  acid,  H.SiF6  ,  precipitates  white,  crystalline  barium  fluo- 
silicate,  BaSiFa  ,  slightly  soluble  in  water  (1-4000),  not  soluble  in  alcohol 
(distinction  from  strontium  and  calcium).  If  an  equal  volume  of  alcohol  be 
added  the  precipitation  is  complete,  sulphuric  acid  not  giving  a  precipitate  in 
the  filtrate  (Fresenius,  Z.,  1890,  29,  143). 

7.  Ignition. — The  volatile  salts  of  barium  as  the  chloride  or  nitrate  impart  a 
yellowish-green  color  to  the  flame  of  the  Bunsen  burner,  appearing  blue  when 
viewed  through  a  green  glass.  The  spectrum  of  barium  is  readily  distinguished 
from  the  spectra  of  other  metals  by  the  green  bands  Ba",  ft  and  y.  Barium 
carbonate  is  very  stable  when  heated,  requiring  a  very  high  heat  to  decompose 
it  into  BaO  and  CO2  . 


STRONTIUM.  §186, 8. 

8.  Detection. — In  the  filtrate  from  the  fourth  group,  barium  is  precipi- 
tated with  strontium  and  calcium  as  the  carbonate  by  ammonium  car- 
bonate. The  white  precipitate  (well  washed)  is  dissolved  in  acetic  acid 
and  the  barium  precipitated  with  K2Cr207  as  BaCr04  which  separates  it 
from  strontium  and  calcium.  The  barium  is  further  identified  by  the 
non-solubility  of  the  chromate  in  acetic  acid,  the  solubility  in  hydrochloric 
acid,  and  precipitation  from  this  solution  by  sulphuric  acid.  It  may  also 
be  confirmed  by  the  color  of  the  flame  with  any  of  the  volatile  salts  (7) 
(not  the  sulphate). 

9.  Estimation. — Barium  is  weighed  as  a  sulphate   (Fresenius  and  Hurtz,  Z. 
angew.,  1896,   253),   carbonate   or   fluosilicate    (BaSiF0).     It   is   separated   from 
strontium  and  calcium:  (1)  "By  digesting  the  mixed  sulphates  at  ordinary  tem- 
peratures for  12  hours  with  ammonium  carbonate.     The  calcium  and  strontium 
are   thus   converted   into   carbonates,    which   are    separated    from   the    barium 
sulphate    by    dissolving    in    hydrochloric    acid.     (2)    By    hydrofluosilicic    acid. 
(3)  By  repeated  precipitation  as  the  chromate  in  an  acetate  solution. 

It  is  separated  from  calcium  by  the  solution  of  the  nitrate  of  the  latter  in 
amyl  alcohol  (§188,  9).  The  hydroxide  and  carbonates  are  also  determined  by 
alkalimetry.  Volumetrically  it  is  precipitated  as  the  chromate,  thoroughly 
washed,  dissolved  in  dilute  HC1  and  the  Crvi  determined  by  H2O2  (Baumann, 
Z.  angew.,  1891,  331). 

10.  Oxidation. — Barium  compounds  are  reduced  to  the  metal  when   heated 
with  Na  or  K   (3).     Ba02   oxidizes  MnCL   to  Mn.,O3    (Spring  and  Lucion,  Bl.t 
1890,  (3),  3,  4). 

§187.  Strontium.    Sr  =  87.63.    Valence  two. 

1.  Properties. — Specific  gravity,  2.4   (Franz,  J.  pr.,   1869,   107,  254).     Melting 
point,   >810°,       850°  ?  (C'ir.  B.  of  S.,   1915),  and  is  not  volatile  when    heated 
to  a  full  red.     It  is  a  "brass-yellow"  metal,  malleable  and  ductile.     It  oxidizes 
rapidly  when  exposed  to  the  air,  and   when  heated  in  the  air  burns,    as    does 
barium,  with  intense  illumination  (Franz,  I.  c.). 

2.  Occurrence. — Strontium    occurs    chiefly    in    strontianite,    SrCO3  ,    and  in 
celestite,  SrSO,  . 

3.  Preparation. — First  isolated  in  1808  by  Davy  by  electrolysis  of  the  hydrox- 
ide (Tran-s.  Royal  $oc.,  345).     It  is  made  by  electrolysis  of  the  chloride  (Frey, 
A.,    1876,    183,    367);    by   heating  a    saturated    solution    of   SrCl,    with    sodium 
amalgam  and  distilling  off  the  mercury  (Franz,  7.  c.);  by  heating  the  oxide  with 
powdered  magnesium  the  metal  is  obtained  mixed  with  MgO  (Winkler,  B.,  1890, 
23,  125). 

4.  Oxides  and  Hydroxides. — Strontium  oxide,  SrO  ,  is  formed  by  igniting  the 
hydroxide,   carbonate    (greater  heat   required    than  with   calcium   carbonate), 
nitrate  and  all  organic  strontium  salts.     The  hydroxide,  Sr(OH)2  ,  is  formed 
by  the  action  of  water  on  the  oxide.     The  peroxide,  SrO2.8H2O  ,  is  made  by  pre- 
cipitating the  hydroxide  with  H2O2;  at  100°  this  loses  water  and  becomes  SrO2  , 
a  white   powder,   melting  at   a  red   heat,   used   in   bleaching   works    (Conroy, 
J.  8oc.  /•»(?..,  1892,  11,  812). 

5.  Solubilities. — a. — Metal. — Strontium    decomposes   water    at    ordinary    tem- 
perature   (Winkler,    I.  c.),   it   is   soluble   in   acids   with   evolution   of   hydrogen, 
ft. — Oxides  and  hydroxides. — The  oxide,  SrO  ,  is  soluble  in  about  100  parts  water 
at  ordinary  temperature,  and  in  about  five  parts  of  boiling  water  forming  the 
hydroxide    (Scheibler,    Neue   Zeitsctirift   fur  Rve&tt&udcer,    1881,    49,    257).     The 
peroxide  is  scarcely  soluble  in  water  or  in  ammonium  hydroxide,  soluble  in 
acids  and  in  ammonium  chloride. 


§187,  6/1.  STRONTIUM.  215 

c. — Salts. — The  chloride  is  slightly  deliquescent;  crystals  of  the  nitrate 
and  acetate  effloresce.  The  chloride  is  soluble,  the  nitrate  insoluble  in 
absolute  alcohol.  The  nitrate  is  insoluble  in  boiling  amyl  alcohol  (§188, 
be).  The  sulphate  is  very  sparingly  soluble  in  water  (1-10,090  at  20.1°) 
(Hollemann,  Z.  pliys.  Ch.,  1893,  12,  131);  yet  sullidently  soluble  to  allow 
its  use  as  a  reagent  to  detect  the  presence  of  traces  of  barium.  Less  soluble 
in  water  containing  ammonium  salts,  sodium  sulphate,  or  sulphuric  acid 
than  in  pure  water;  quite  appreciably  soluble  in  HC1  or  HNO.<  ;  insoluble 
in  alcohol.  Strontium  nuosilicate  is  soluble  in  water  (distinction  from 
barium).  The  chromate  is  soluble  in  831.8  parts  water  at  15°  (Fresenius, 
Z.,  1890,  29,  419);  soluble  in  many  acids  including  chromic  acid;  and  more 
soluble  in  water  containing  ammonium  salts  than  in  pure  water. 

6.  Reactions,  a. — The  fixed  alkalis  precipitate  strontium  salts  when 
not  too  dilute,  as  the  hydroxide,  Sr(OH)., ,  less  soluble  than  the  barium 
hydroxide.  No  precipitate  with  ammonium  hydroxide.  The  alkali  car- 
bonates precipitate  solutions  of  strontium  salts  as  the  carbonate.  Stron- 
tium sulphate  is  completely  transposed  on  boiling  with  a  fixed  alkali- car- 
bonate (distinction  from  barium,  §188,  6a  footnote). 

6. — Oxalic  acid  and  oxalates  precipitate  strontium  oxalate,  insoluble  in 
water,  soluble  in  hydrochloric  acid  (Souchay  and  Lenssen,  A.,  1857,  102,  35). 
c. — The  solubility  of  strontium  salts  is  diminished  by  the  presence  of  con- 
centrated nitric  acid,  but  less  so  than  barium  salts,  d. — In  deportment  with 
phosphates,  strontium  is  not  to  be  distinguished  from  barium. 

e. — See  6e,  §§186  and  188.  Sulphuric  acid  and  sulphates  (including 
CaS04)  precipitate  solutions  of  strontium  salts  as  the  sulphate,  unless 
the  solution  is  diluted  beyond  the  limit  of  the  solubility  of  the  precipitate 
(5c).  A  solution  of  strontium  sulphate  is  used  to  detect  the  presence  of 
traces  of  barium  (distinction  from  strontium  and  calcium).  In  dilute 
solutions  the  precipitate  of  strontium  sulphate  forms  very  slowly,  aided 
by  boiling  or  by  the  presence  of  alcohol,  prevented  by  the  presence  of 
hydrochloric  or  nitric  acids  (5c).  It  is  almost  insoluble  in  a  solution  of 
ammonium  sulphate  (separation  from  calcium). 

f. — The  hulides  of  strontium  are  all  soluble  in  water  and  have  no  application 
in  the  analysis  of  strontium  salts.  Strong-  hydrochloric  acid  dissolves  stron- 
tium sulphate,  but  in  general  diminishes  the  solubility  of  strontium  salts  in 
water,  g. — Neutral  solutions  of  arsenites  do  not  precipitate  strontium  salts. 
The  addition  of  ammonium  hydroxide  causes  a  precipitation  of  a  portion  of  the 
strontium.  Arsenate  of  strontium  resembles  the  corresponding  barium  salt. 
Alkaline  arsenates  do  not  precipitate  strontium  from  solution  of  the  sulphate 
(distinction  from  calcium,  §188,  60). 

h. — Normal  chromates  precipitate  strontium  chromate  from  solutions 
not  too  dilute  (5c),  soluble  in  acids.  In  absence  of  barium,  strontium 
may  be  separated  from  calcium  by  adding  to  the  nearly  neutral  solutions 
a  solution  of  K2Cr04  plus  one-third  volume  of  alcohol.  The  calcium 


216  CALCIUM.  §187, 6i. 

chromate  is  about  100  times  as  soluble  as  the  strontium  chromate  (Fre- 
senius  and  Rubbert.  Z.,  1891,  30,  672).  No  precipitate  is  formed  with 
potassium  bichromate  (separation  from  barium). 

i. — Fluosilicic  acid  does  not  precipitate  strontium  salts  even  from  quite 
concentrated  solutions,  as  the  strontium  fluosilicate  is  fairly  soluble  in  cold 
water  and  more  so  in  the  presence  of  hydrochloric  acid  (Fresenius,  Z.,  1890, 
29,  143). 

7.  Ignition. — Volatile  strontium  compounds  color  the  flame  rriuwtn.  In  pres- 
ence of  barium  the  crimson  color  appears  at  the  moment  when  the  substance 
(moistened  with  hydrochloric  acid,  if  a  non-volatile  compound)  is  first  brought 
into  the  flame.  The  paler,  yellowish-red  flame  of  calcium  is  liable  to  be  mis- 
taken for  the  strontium  flame.  The  spectrum  of  strontium  is  characterized 
by  eight  bright  bands;  namely,  six  red,  one  orange  and  one  blue.  The  orange 
line  Sr  «,  at  the  red  end  of  the  spectrum;  the  two  red  lines,  Sr  ft  and  Sr  y, 
and  the  blue  line,  Sr  6 ,  are  the  most  important. 

8.  Detection. —  Strontium  is  precipitated  with  barium  and  calcium  from 
the  filtrate  of  the  fourth  group  by  ammonium  carbonate.  The  well  washed 
precipitate  of  the  carbonates  is  dissolved  in  acetic  acid  and  the  barium 
removed  by  K2Cr207 .  The  strontium  and  calcium  are  separated  from  the 
excess  of  chromate  by  reprecipitation  with  (NH4)2C03 .  The  precipitate  is 
again  dissolved  in  HC2H302  and  from  a  portion  of  the  solution  the  stron- 
tium is  detected  by  a  solution  of  CaS04  (6e).  The  flame  test  (7)  is  of  value 
in  the  identification  of  strontium. 

9.  Estimation. — Strontium  is  weighed  as  a  sulphate  or  a  carbonate.  The 
hydroxide  and  carbonate  may  be  determined  by  alkalimetry.  It  is  separated 
from  calcium:  (1)  By  the  insolubility  of  its  sulphate  in  ammonium  sulphate. 
(2)  By  boiling  the  nitrates  with  amyl  alcohol  (§188,  9).  (3)  By  treating  the 
nitrates  with  equal  volume  of  absolute  alcohol,  and  ether  (§188,  9).  For 
separation  from  barium  see  §186,  9. 

§188.  Calcium.    Ca  =  40.07.    Valence  two. 

1.  Properties.—  Specific  gravity,  1.6  to  1.8  (Caron,  C.  r.,  1860,  50,  547).     Melting 
point,  810°  (Cir.  B.  of  S.,  1915).     A  white  metal  having  very  much  the  appearance 
of  aluminum,  is  neither  ductile  nor  malleable   (Frey,   A.,   1876,   183,   367).     In 
dry  air  it  is  quite  stable,  in  moist  air  it  burns  with  incandescence,  as  it  does  also 
with  the  halogens.     It  dissolves  in  mercury,  forming  an  amalgam. 

2.  O3cunence. — Found   in   the    mineral   kingdom  *as   a   carbonate  in  marble, 
limestone,   chalk  and  aragonite;    as  a  sulphate  in  gypsum,   selenite,   alabaster, 
etc.;  as  a  fluoride  in  fluor-spar;  as  a  phosphate  in  apatite,  phosphorite,  etc. 
It  is  found  as  a  phosphate  in  bones;  in  egg-shells  and  oyster-shells,  as  a  car- 
bonate.    It  is  found  in  nearly  all  spring  and  river  waters'. 

3.  Preparation. —  (1)  By  ignition  of  the  iodide  with  sodium  in  closed  retorts 
(Dumas,  C.  r.,  1858,  47,  575).     (2)   By  fusion  of  a  mixture  of  300  parts  fused 
CaCL  ,  400  parts  granulated  zinc  and  100  parts  Na  until  zinc  vapor  is  given 
off.     From  the  CaZn  alloy  thus  obtained  the  zinc  is  removed  by  distillation  in 
a  graphite  crucible  (Caron,  I.  c.).    (3)  By  electrolysis  of  the  chloride  (Frey,  I.  c.). 
(4)  By  reducing  the  oxide,  hydroxide  or  carbonate  with  magnesium  (Winkler, 
B.,  1890,  23,  122  and  2642). 

4.  Oxides  and  Hydroxides.— The  oxide,  CaO  ,  is  a  strong  base,  non-fusible, 
non-volatile;  it  is  formed  by  oxidation  of  the  metal  in  air;  by  ignition  of  the 


§188,  5c.  CALCIUM.  217 

hydroxide,  the  carbonate  (limestone),  nitrate,  and  all  organic  calcium  salts. 
The  corresponding1  hydroxide,  Ca(OH)2  (slaked  lime),  is  made  by  treating  the 
oxide  with  •water.  Its  usefulness  when  combined  with  sand,  making-  mortar, 
is  too  well  known  to  need  any  description  here.  The  peroxide,  CaO,.sH,0  ,  is 
made  by  adding-  hydrogen  peroxide  or  sodium  peroxide  to  the  hydroxide: 
Ca(OH)!,  +  H2O,  ==  Ca02  +  2H,O  (Conroy,  J.  Nor.  Intl.,  1892,  11,  80S).'  Drying 
at  130°  removes  all  Ilie  watei1,  leaving-  a  \vhi1c  powder,  CaO,  ,  which  at  a  red 
heat  loses  half  its  oxygen  (Schoene,  A.,  3877,  192,  257).  It  cannot  be  made  by 
heating  the  oxide  in  oxygen  or  with  potassium  chlorate  (§186,  4). 

r>.  Solubilities. — a. — Metal. — Calcium  is  soluble  in  acids  with  evolution  of 
hydrogen;  it  decomposes  water,  evolving  hydrogen  and  forming  Ca(OH),  . 

&. — Oxide  and  hydroxide. — CaO  combines  with  dilute  acids  forming  cor- 
responding salts,  it  absorbs  C02  from  the  air  becoming  CaCO,  .*  In  moist 
air  it  becomes  Ca(OH)2  ,  the  reaction  taking  place  rapidly  and  with  increase 
of  volume  and  generation  of  much  heat  in  presence  of  abundance  of 
water.  The  hydroxide,  Ca(OH)2 ,  is  soluble  in  acids,  being  capable  of 
titration  with  standard  acids.  It  is  much  less  soluble  in  water  than 
barium  or  strontium  hydroxides  (Lamy,  C.  r.,  1878,  86,  333);  in  806  parts 
at  19.5°  (Paresi  and  Eotondi,  7?.,  1874,  7,  817);  and  in  1712  parts  at  100° 
(Lamy,  /.  c.).  The  solubility  decreases  with  increase  of  temperature.  In 
saturated  solutions  one  part  of  the  oxide  is  found  in  744  parts  of  water 
at  15°  (Lamy,  I.  c.).  A  clear  solution  of  the  hydroxide  in  water  is  lime 
water  (absorbs  C02  forming  CaCO.,),  the  hydroxide  in  suspension  to  a 
greater  or  less  creamy  consistency  is  milk  of  lime. 

c. — Salts. — The  chloride,  bromide,  iodide,  nitrate,  and  chlorate  are 
deliquescent;  the  acetate  is  efflorescent. 

The  carbonate,  oxalate,  and  phosphate  are  insoluble  in  water.  The 
chloride,  iodide,  and  nitrate  are  soluble  in  alcohol.  The  nitrate  is  soluble 
in  1.87  parts  of  equal  volumes  of  ether  and  alcohol  (Fresenius,  Z.,  1893, 
32,  191);  readily  soluble  in  "boiling  amyl  alcohol  (Browning,  Am.  S.,  1892, 
143,  53  and  314)  (separation  from  barium  and  strontium).  The  carbonate 
is  soluble  in  water  saturated  with  carbonic  acid  (as  also  are  barium,  stron- 
tium, and  magnesium  carbonates),  giving  hardness  to  water.  The  oxalate 
is  insoluble  in  acetic  acid,  soluble  in  hydrochloric  and  nitric  acids.  The 
sulphate  is  soluble  in  about  500  parts  of  water  f  at  ordinary  temperature, 
the  solubility  not  varying  much  in  hot  water  until  above  100°  when  the 
solubility  rapidly  decreases.  Its  solubility  in  most  alkali  salts  is  greater 
than  in  pure  water.  Ammonium  sulphate  (1-4)  requires  287  parts  for  the 
solution  of  one  part  of  CaS04  (Fresenius,  Z.,  1891,  30,  593)  (separation 
from  Ba  and  Sr).  Eeadily  soluble  in  a  solution  of  Na2S203  (separation 
from  barium  sulphate)  (Diehl,  J.  pr.,  1860,  79,  430).  It  is  soluble  in  60 
parts  hydrochloric  acid,  6.12  per  cent  at  25°,  and  in  21  parts  of  the  same 

*  Dry  CaO  does  not  absorb  dry  CO2  or  SO,  below  350°.  (Veley,  J.  C.,  1893,  63,  82K 
t  Goldhammer,  C.  C.,  1888,  708;  Droeze,  B.,  18T7,  I  o.  330;  lioisbuudran,  A.  C7).,  1874.  (5),  3,477 
Kohlrausch  and  Rose,  Z.  phys.  C?i.,  1893, 12, 241 ;  Raupenstrauch,  M.,  1885,  6, 663). 


218  CALCIUM.  §188,  6fl. 

acid  at  103°  (Lunge,  J.  Soc.  Ind.,  1895  14,  31).  The  chromate  is  soluble 
in  214.3  parts  water  at  14°  (Siewert,  J.,  1862,  149);  in  dilute  alcohol  it  is 
rather  more  soluble  (Fresenius,  /.  c.,  page  672);  very  readily  soluble  in 
acids  including  chromic  acid. 

6.  Reactions,  a.—  The  fixed  alkali  hydroxides  precipitate  solutions  of 
calcium  salts  not  having  a  degree  of  dilution  beyond  the  solubility  of  the 
calcium  hydroxide  formed  (5&),  i.  e.  potassium  hydroxide  will  form  a 
precipitate  with  calcium  sulphate  since  the  sulphate  requires  less  water 
for  its  solution  than  the  hydroxide  (56  and  c);  also  the  calcium  hydroxide 
is  less  soluble  in  the  alkaline  solution  than  in  pure  water.  Ammonium 
hydroxide  does  not  precipitate  calcium  salts.  The  alkali  carbonates  pre- 
cipitate calcium  carbonate,  CaC03 ,  insoluble  in  water  free  from  carbon 
dioxide,  decomposed  by  acids.  Calcium  sulphate  is  completely  trans- 
posed upon  digestion  with  an  alkali  carbonate  *  (distinction  from  barium). 
Calcium  hydroxide,  Ca(OH)2 ,  is  used  as  a  reagent  for  the  detection  of 
carbon  dioxide  (56  and  §228,  8). 

6. — Alkali  oxalates,  as  (NH4)2C204 ,  precipitate  calcium  oxalate,  CaC204 , 
from  even  dilute  solutions  of  calcium  salts.  The  precipitate  is  scarcely  at 
all  soluble  in  acetic  or  oxalic  acids  (separation  of  oxalic  from  phosphoric 
acid  (§315),  but  is  soluble  in  hydrochloric  and  nitric  acids.  The  pre- 
cipitation is  hastened  by  presence  of  ammonium  hydroxide.  Formed 
slowly,  from  very  dilute  solutions,  the  precipitate  is  crystalline,  octahedral. 
If  Sr  or  Ba  are  possibly  present  in  the  solution  to  be  tested  (qualitatively), 
an  alkali  sulphate  must  first  be  added,  and  after  digesting  a  few  minutes, 
if  a  precipitate  appears,  SrS04  ,  BaS04 ,  or,  if  the  solution  was  concentrated, 
perhaps  CaS04 ,  it  is  filtered  out,  and  the  oxalate  then  added  to  the  filtrate. 
If  a  mixture  of  the  salts  of  barium,  strontium,  and  calcium  in  neutral  or 
alkaline  solution  be  treated  with  a  mixture  of  (NH4)2S04  and  (NH4)2C204 , 
the  barium  and  strontium  are  precipitated  as  sulphates  and  the  calcium  as 
the  oxalate;  separated  from  the  barium  and  strontium  on  addition  of 
hydrochloric  acid  (Sidersky,  Z.,  1883,  22,  10;  Bozomoletz,  B.,  1884,  17, 
1058).  A  solution  of  calcium  chloride  is  used  as  a  reagent  for  the  detec- 
tion of  oxalic  acid  (§227,  8). 

In  solutions  of  calcium  salts  containing1  a  strong1  excess  of  ammonium 
chloride,  potassium  ferrocyanide  precipitates  the  calcium  (distinction  from 
barium  and  strontium)  (Baubigny,  Bl.,  1895,  (3),  13,  326). 

*  Here  experiment  shows  that  for  equilibrium  the  SO4  ions  must  be  present  in  solution  in  large 
excess  of  CO3  ions.  With  strontium  also  an  excess  of  SO4  ions  is  required,  although  not  so 
great  as  in  the  case  of  calcium.  For  barium,  however,  equilibrium  demands  that  the  concen- 
tration of  CO3  ions  exceed  that  of  SO4.  This  condition  is  already  fulfilled  when  an  alkali  car- 
bonate is  added  to  BaSO4  and  therefore  no  change  takes  place  in  this  case,  while  in  the  others 
the  sulphate  is  transformed  into  carbonate.  It  is  important  to  notice  that  the  relative  or  ab- 
solute quantities  of  solid  carbonate  and  sulphate  present  do  not  affect  the  equilibrium,  which 
is  determined  solely  by  the  substances  in  solution  (§57,  6e,  footnote). 


§188, 9.  CALCIUM. 

c. — See  5^.  d. — By  the  action  of  alkali  phosphates,  solutions  of  calcium  are 
not  distinguished  from  solutions  of  Iwriiim  or  strontium. 

e. — Pure  sodium  sulphide,  NaL.S  ,  gives  an  abundant  precipitate  with  calcium 
salts;  even  with  CaSO4  .  The  precipitate  is  Ca(OH),:  CaCL  +  2Na2S  +  2H.O  = 
Ca(OH)2  +  2NaCl  +  2NaHS  .  The  acid  sulphide,  NaHS  ,  does  not  precipitate 
calcium  salts  (Pelouze,  A.  Cli.,  1866,  (4),  7,  172).  Alkali  sulphites  precipitate 
calcium  sulphite,  nearly  insoluble  in  water,  soluble  in  hydrochloric,  nitric  or 
sulphurous  acid;  barium  and  strontium  salts  act  similarly. 

Sulphuric  acid  and  soluble  sulphates  precipitate  calcium  salts  as  CaS04 , 
distinguished  from  barium  by  its  solubility  in  water  and  in  hydrochloric 
acid;  from  barium  and  strontium  by  its  solubility  in  ammonium  sulphate 
(5c).  A  water  solution  of  calcium  sulphate  is  used  to  detect  strontium 
after  barium  has  been  removed  as  a  chromate.  Obviously  a  solution  of 
strontium  sulphate  will  not  precipitate  calcium  salts. 

/. — Calcium  chloride,  fused,  is  much  used  as  a  drying  agent  for  solids,  liquids 
and  gases.  Chlorinated  lime,  or  bleaching  powder,  CaCJUO  (Kingzett,  J.  C., 
1875,  28,  404),  is  much  used  as  a  bleaching  agent  and  as  a  disinfectant,  g. — 
Neutral  or  ammoniacal  solutions  of  arsenites  form  a  precipitate  with  calcium 
salts  (distinction  from  barium).  A  solution  of  calcium  salts  including  solu- 
tions of  calcium  sulphate  in  ammoniacal  solution  is  precipitated  by  arsenic 
acid  as  CaNH4As04  (distinction  from  strontium  after  the  addition  of  sulphuric 
acid)  (Bloxam,  C.  N.,  1886,  54,  16). 

h. — Normal  chromates,  as  K2Cr04  ,  precipitate  solutions  of  calcium  salts  as 
calcium  chromate,  CaCr04  ,  yellow,  provided  the  solution  be  not  too  dilute  (5c). 
The  precipitate  is  readily  soluble  in  acids  and  is  not  formed  with  acid  chro- 
mates as  K2Cr2OT  (separation  from  barium),  i. — Fluosilicic  acid  does  not 
precipitate  calcium  salts  even  in  the  presence  of  equal  parts  of  alcohol  (separa- 
tion from  barium). 

7.  Ignition. — Calcium  sulphate,  CaS04.2H20  ,  gypsum  ,  loses  its  water  of 
crystallization  at  80°  and  becomes  the  anhydrous  sulphate,  CaS04  ,  plaster  of 
Paris;  which  on  being  moistened  forms  the  crystalline  CaS04.2H2O  ,  expands 
and  "  sets."  Calcium  carbonate,  limestone,  when  heated  (burned)  loses  carbon 
dioxide  and  becomes  lime,  CaO  . 

Compounds  of  calcium,  preferably  the  chloride,  render  the  flame  yellowish 
red.  The  presence  of  strontium  or  barium  obscures  this  reaction,  but  a  mixture 
containing  calcium  and  barium,  moistened  with  hydrochloric  acid,  gives  the 
calcium  color  on  its  first  introduction  to  the  flame.  The  spectrum  of  calcium 
is  distinguished  by  the  bright  green  line,  Ca  /?,  and  the  intensely  bright 
orange  line,  Ca  a,  near  the  red  end  of  the  spectrum. 

8.  Detection. — Calcium  is  separated  in  analysis  from  the  metals  of  the 
other  groups  and  from  barium,  with  strontium,  as  described  at  §187,  8. 
A  portion  of  the  solution  of  strontium  and  calcium  acetate  is  boiled  with 
potassium  sulphate;  after  standing  for  some  time  (ten  minutes),  the  filtrate 
is  tested  with  ammonium  oxalate.  A  white  precipitate  insoluble  in  the 
acetic  acid  present,  but  soluble  in  hydrochloric  acid  is  evidence  of  the 
presence  of  calcium.  The  flame  test  (7)  is  confirmatory. 

9.  Estimation. — Calcium  is  weighed  as  an  oxide,  carbonate,  or  sulphate.  The 
carbonate  is  obtained  by  precipitating  as  oxalate,  and  gently  igniting  the  dried 
precipitate;  higher  ignition  changes  the  carbonate  to  the  oxide.  The  sulphate 
is  precipitated  in  a  mixture  of  two  parts  of  alcohol  to  one  of  the  solution.  The 
hydroxide  and  carbonate  may  be  determined  by  alkalimetry.  Calcium  may  be 
separated  from  barium  and  strontium  by  the  solution  of  its  nitrate  in  amyl 


MAGXE8WM.  §189,  1. 

alcohol  (5c).  The  best  method  of  separation  from  strontium  is  to  treat  the 
nitrates  with  a  mixture  of  equal  volumes  of  alcohol  and  ether.  The  calcium 
nitrate  dissolves,  but  not  more  than  one  part  in  60,000  of  the  strontium  is 
found  in  the  solution  (§195).*  In  the  presence  of  iron,  aluminum  and  phos- 
phoric acid,  calcium  is  best  precipitated  as  an  oxalate  in  the  presence  of  citric 
acid  (Passon,  Z.  angew.,  1898,  776).  See  also  9,  §186  and  §187. 

§189.  Magnesium.    Mg  =  24.32.     Valence  two. 

1.  Properties.—  Specific  gravity,    1.75  (Deville  and    Caron,   A.  Ch.,  1863,   (3), 
67,  346);   melting  point,  651°  (Cir.  B.  of  S.,  1915).     A  white,  hard,  malleable  and 
ductile    metal;     not   acted   upon   by   water   or   alkalis   at    ordinary    temperature 
and  only  slightly  at  100°  (Ballo,  B.,  1883,  16,  694).     When  heated  in  air  or  in 
oxygen  it  burns  with  incandescence  to  MgO  .     It  combines  directly  when  heated 
in  contact  with  N  ,  P  ,  As  ,  S  and  Cl .     It  forms  alloys  with  Hg  and  Sn,    forming 
compounds  which  decompose  water. 

2.  Occurrence. — Magnesite,' MgC03 ;  dolomite,  CaMg(C03)2;  brucite.  Mg(OH)2; 
epsom  salts,  MgSO4.7H20;  and  combined  with  other  metals  in  a  great  variety 
of  minerals. 

3.  Preparation. —  (1)  By  electrolysis  of  the  chloride  or  sulphate  (Bunsen,  A., 
1852,    82,    137).     (2)    By   ignition    of   the    chloride    with    sodium    or    potassium 
(Wohler,   A.,   1857,    101,   562).     (3)    Mg2Fe(CN)6    is   ignited   with   Na2C03  ,    and 
this  product  ignited  with  zinc  (Lanterbronn,  German  Patent  No.  39,915). 

4.  Oxide  and  Hydroxide. — Only  one   oxide   of  magnesium,  MgO  ,   is  known 
with  certainty.     Formed  by  burning  the  metal  in  the  air,  and  by  action   of 
heat  upon  the  hydroxide,  carbonate,  nitrate,  sulphate,  oxalate  and  other  mag- 
nesium salts  decomposed  by  heat.     The  corresponding  hydroxide,  Mg(OH)2  , 
is  formed  by  precipitating  magnesium  salts  with  the  fixed  alkalis. 

5.  Solubilities. — a. — Metal. — Magnesium  is   soluble  in  acids   including 
carbonic  acid,  evolving  hydrogen:    Mg  -f  C02  +  H20  —  MgC03  +  ^ 
(Ballo,  B.,  1882,  15,  3003):  it  is  also  attacked  b}  fhe  acid  alkali  carbonate?, 
as  NaHC03,  to  form  MgC03 ,  Na2C03  and  H  (Ballo,  I.  c.).     Soluble  ID 
ammonium  salts:  Mg  +  3NH4C1  ==  NH4MgCL  +  2NH3  +  H2 .     With 
the  halogens  it  acts  tardily  (Wanklyn  and  Chapman,  J.  C.,  18G6,  19,  141). 
&. — Oxide  and  hydroxide. — Insoluble  in  water,  soluble  in  acids.     Mg(OH)2 
is  soluble  in  111,111  parts  of  water  at  18°  (Kohlrausch  and  Eose,  Zeit. 
phys.   Ch.,   1893,   12,   241).     In  contact  with  water  the  oxide   is   slowly 
changed  to  the  hydroxide,  Mg(OH)2 ,  and  absorbs  C02  from  the  air.     Sol- 
uble   in    ammonium    salts:  *    Mg(OH)2    +    3NH4C1    =    NH4MgCL     + 
2NH4OH  .     c. — Salts. — The   chloride,   bromide,  iodide,   chlorate,   nitrate, 
and  acetate  (4  aq)  are  deliquescent',  the  sulphate  (7  aq)  slightly  efflorescent. 
The  carbonate,  phosphate,  borate,  arsenite,  and  arsenate  are  insoluble  in 
water;  the  sulphite,  oxalate,  and  chromate  soluble;  the  tartrate  sparingly 
soluble.     The  carbonate  is  soluble;  the  phosphate,  arsenite,  and  arsenate 
are  insoluble  in  excess  of  ammonium  salts. 

6.  Reactions,     a. — The  fixed  alkali  hydroxides  and  the  hydroxides  of 
barium,  strontium  and  calcium  precipitate  magnesium  hydroxide,  Mg(OH)3; 

*  The  conditions  here  are  the  same  as  in  the  case  of  Mn(OH)2,  §134,  6a,  footnote. 


§189,  6/.  MAGNESIUM.  221 

white,  gelatinous,  from  solutions  of  magnesium  salts  ;  insoluble  in  excess 

of  the  reagent  but  readily  soluble  in  ammonium  salts,  the  magnesium  pass- 

++    —  +     -        +-f 

ing  into  the  negative  ion:  Mg(OH)2+  4NH4Cl=-(NH4_)2MgCl4+2NH4OH  . 
With  ammonium  hydroxide  but  half  of  the  magnesium  is  precipitated, 

the  remainder  being  held  in  solution  in  the  acid   ion  by  the  ammonium 

++  -  +      -        ++  —  +  + 

salt   formed    in    the   reaction:     2Mg  S04  ~f  2NH4OH  —  Mg(OH)2  -f-  (NHt)2 

Mg(S04)2  (Rheineck,  DingL,  1871,  202,  268).  The  fixed  alkali 
carbonates  precipitate  basic  magnesium  carbonate,  Mg,  (OH)2(C03)3  , 
variable  to  Mg,(OH),(CO:{)4  :  4MgS04  +  INa.CO,  -f  H,0  =  =  Mg4- 
(OH),(C03)3  +  Na,S04  +  C02.  If  the  above  reaction  takes  place  in 
the  cold  the  carbon  dioxide  combines  with  a  portion  of  the  magnesium 
carbonate  to  form  a  soluble  acid  magnesium  carbonate:  5MgS04  -(- 
5Na2C03  +  2H20  ==  Hg4(OH)2(C03)3  +  MgH2(CO,)2  +  5Na2S04  .  On 
boiling,  the  acid  carbonate  is  decomposed  with  escape  of  CO,  .  Ammonium 
carbonate  docs  not  precipitate  magnesium  salts,  as  a  soluble  double  salt  is 
at  once  formed.  Acid  fixed  alkali  carbonates,  as  NaHC03  ,  do  not  precipi- 
tate magnesium  salts  in  the  cold;  but  upon  boiling,  C02  is  evolved  and  the 
carbonate  is  precipitated  (Engel,  A.  Ch.,  1886,  (6),  7,  260). 

1).  —  Soluble  oxalates  do  not  precipitate  solutions  of  magnesium  salts,  as  they 
form  soluble  double  oxalates.  If  to  the  solution  of  double  oxalates,  preferably 
magnesium  ammonium  oxalate,  an  equal  volume  of  80  per  cent  acetic  acid  be 
added,  the  magnesium  is  precipitated  as  the  oxalate  (separation  from  potas- 
sium or  sodium  (Classen,  Z.,  1879,  18,  373). 

d.  —  Alkali  phosphates  —  as  Na2HP04  —  precipitate  magnesium  phosphate, 
MgHP04  ,  if  the  solution  be  not  very  dilute.     But  even  in  very  dilute 
solutions,  by  the  further  addition  of  ammonium  hydroxide  (and  NH4C1)3 
a  crystalline  precipitate  is  slowly  formed,  magnesium  ammonium  phosphate 
—  Mg-ini4P04  .     Stirring  with  a  glass  rod  against  the  side  of  the  test-tube 
•promotes  the  precipitation.     The  addition  of  ammonium  chloride,   in  this 
test,  prevents  formation  of  any  precipitate  of  magnesium  hydroxide  (5&). 
The  precipitate  dissolves  in  13,497  parts  of  water  at  23°   (Ebermayer, 
J.  pr.,   1853,   60,   41);  almost   absolutely  insoluble  in  water  containing 
ammonium  hydroxide  and  ammonium  chloride  (Kubel,  Z.9  1869,  8,  125). 

e.  —  Magnesium  xulphide  is  decomposed  by  water,   and   magnesium  salts  are 
not  precipitated  by  hydrosulphuric  acid  or  ammonium  sulphide;  but  Mg-0  -f- 
H2O    (1-10)    absorbs   H2S  ,   forming-   in   solution    MgH^S.  ,    which   readily   gives 
off  HJ3  upon  boiling  (a  very  satisfactory  method  of  preparing  H.S  absolutely 
arsenic   free)    (i)ivers  and   Shmidzu,  J.   C.,   1884,  45,   699).     Normal  sodium  or 
potassium  sulphide  precipitates  solutions  of  magnesium  salts  as  the  hydroxide 
with  formation  of  an  acid  alkali  sulphide:  MgSO',  +  2Na.,S  -f  2H,O  =  Mg(OH), 
+  Na,S04   +  2NaHS   (IVIou/c.  .-I.  CIi..   1866,   (4),  7,  172).     Sulphuric  acid  and 
soluble  sulphates  do  not  precipitate  solutions  of  magnesium  salts  (distinction 
from  Ba  ,  Sr  and  Ca). 

f.  —  Magnesium  chloride,  in  solution,  evaporated  on  the  water  bath   evolves 


CALIFORNIA   COLLEi* 


222  MAGNESIUM.  §189,  60. 

hydrochloric  ac:d  (7).     g. — Soluble   arsenates  precipitate  magnesium  salts  in  de- 
portment similar  to  the  corresponding  phosphates. 

7.  Ignition. — Magnesium  ammonium  phosphate  when  ignited  loses  ammonia 
and  water,  and  becomes  the  pyrophosphate:  SMg-NHiPC^  =  Mg2P207  +  HaO  + 
2NH3  .     The  carbonate  loses  CO2  and  becomes  MgO  .     In  dry  air  magnesium 
chloride  may  be  ignited  without  decomposition,  but  in  the  presence  of  steam 
MgO  and  HC1  are  formed:  MgCl2  -f-  H,O  =  MgO  +  2HC1;  a  technical  method 
for  preparing  HC1  (Heumann,  A.,  1877,  184,  227). 

8.  Detection. — If  sufficient  ammonium  salts  have  been  used,  the  mag- 
nesium will  be  in  the  filtrate  from  the  precipitated  carbonates  of  barium, 
strontium  and  calcium.     From  a  portion  of  this  filtrate  the  magnesium  is 
precipitated  as  the  white  magnesium  ammonium-phosphate,  MgNH4P04 , 
by  Na2HP04 . 

9.  Estimation. — After  removal  of  other  non-alkali  metals,  magnesium  is  pre- 
cipitated as  MgNH4PO4  ,  then  changed  by  ignition  to  Mg2P2O7  (magnesium 
pyrophosphate)  and  weighed  as  such.  Separated  as  MgCL  from  KC1  and  NaCl 
by  solution  in  amyl  alcohol,  evaporated  with  H2S04  and  weighed  as  MgS04 
(Riggs,  Am.  $.,  1892,  44,  103).  It  is  estimated  volumetrically  by  precipitation 
as  MgNH4PO4  ,  drying  at  about  50°  until  all  free  NH4OH  is  removed.  An 
excess  of  standard  acid  is  then  added  and  at  once  titrated  back  with  standard 
fixed  alkali,  using  methyl  orange  as  an  indicator  (Handy,  J.  Am.  Soc.,  1900,  22, 
31). 

10.  Oxidation. — Magnesium  is  a  powerful  reducer;  ignited  with  the 
oxides  or  carbonates  of  the  following  elements  magnesium  oxide  is  formed 
and  the  corresponding  element  is  liberated :  Ag ,  Hg ,  Pt ,  Sn  *,  B ,  Al , 
Th,  CJ,  Si,  Pb,  PJ,  As,  Sb,  Bi,  Cr,  Mo,  Mn,  Fe,  Co,  Ni,  Cu, 
Cd ,  Zn  ,  Gl ,  Ba  ,  Sr ,  Ca  ,  Rb  ,  K  ,  Na ,  and  Li .  In  some  cases  the  reaction 
takes  place  with  explosive  violence.  From  their  corresponding  salts  in^ 
neutral  solution  Mg  precipitates  Se ,  Te ,  As ,  Sb  ,  Bi ,  Sn  ,  Zn  f  ,  Cd ,  Pb  , 
Tl ,  Th ,  Cu ,  Ag ,  Mn  f,  Fe  f,  Co ,  Ni ,  Au ,  Pt ,  and  Pd  (Scheibler,  B., 
1870,  3,  295;  Villiers  and  Borg,  C.  r.,  1893,  116,  1524). 

*  Winkler,  B.,  1890,  23, 44, 130  and  772 ;  1891,  24,  888. 

t  Kern,  C.  N ,  1876,  33, 112  and  236. 

t  Seubert  and  Schmidt,  A.,  1892,  267,  818. 


190 


ANALYSIS   OF  THE   CALCIUM  GROUP. 


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224  btRtiCTWNS  FOR  AXAL78TS  WITH  NOTES.  §191, 

DIRECTIONS  FOR  ANALYSIS  OF  THE  METALS  OF  THE  CALCIUM  GROUP. 
(THE  ALKALIXE  EARTHS.) 

§191.  Manipulation. — To  the  filtrate  from  the  fourth  group  in  which 
H2S  (§192,  1)  gives  no  precipitate  (§138)  add  NH4OH  and  ammonium 
carbonate  as  long  as  a  precipitate  is  formed:  BaCl2  -f-  (NH4).,C03  =  BaCOs 
+  2NH4C1 .  Digest  with  warming,  filter  and  wash.  The  filtrate  should 
be  tested  again  with  ammonium  carbonate  and  if  no  precipitate  is  formed 
it  is  set  aside  to  be  tested  for  magnesium  and  the  alkali  metals  (§§193 
and  211). 

The  well  washed  white  precipitate  is  dissolved  in  acetic  acid,  using  as 
little  as  possible :  SrCO,  +  2HC2H302  ==  Sr(C2H,02)2  +  C02  +  H20  . 

To  a  small  portion  of  the  acetic  acid  solution  add  a  drop  of  K.,Cr,07  ; 
if  a  precipitate—  BaCr04—  is  obtained,  the  K2Cr207  must  be  added  to  the 
whole  solution:  2Ba(C2H,02)  +  K2Cr20T  +  ELO  =  2BaCr04  +  2X0,11,0, 
+  2HC2H302 .  Filter,  wash  the  precipitate  and  dissolve  it  in  HC1.  Test 
a  portion  in  the  flame  and  precipitate  the  barium  in  the  remainder  as 
barium  sulphate,  with  a  drop  of  sulphuric  acid. 

To  the  filtrate  from  the  barium  chromate  add  NH4OH  and  (NH4)2C03 , 
warm,  filter,  and  wash.  Dissolve  the  white  precipitates  of  SrC03  and 
CaC03  in  acetic  acid  and  divide  the  solution  into  two  portions. 

Portion  1. — For  Strontium. — With  a  platinum  wire  obtain  the  flame 
test,  crimson  for  strontium;  calcium  interferes  (7,  §§187,  188  and  205). 
Add  a  solution  of  calcium  sulphate  and  boil;  set  aside  for  about  ten  min- 
utes. A  precipitate — Sr£04 — indicates  strontium.  This  SrS04  may  be 
moistened  with  HC1  and  the  crimson  flame  test  obtained. 

Portion  2. — For  Calcium. — Add  a  solution  of  ammonium  (1-4)  sulphate, 
boil,  and  set  aside  for  ten  minutes.  Filter  (to  remove  any  strontium  that 
may  be  present;  also  a  portion  of  the  calcium  may  be  precipitated,  §188,  6e.) 
and  add  ammonium  oxalate  to  the  filtrate.  Dissolve  the  precipitate  in 
HC1,  A  white  precipitate — CaC204 — insoluble  in  acetic  acid  by  its  forma- 
tion in  that  solution,  and  soluble  in  HC1  is  proof  of  the  presence  of  calcium. 

§192.  Notes. — 1.  Considerable  amounts  of  the  metals  of  this  group,  especially 
barium  and  strontium,  may  be  precipitated  with  the  second  group  on  account 
of  the  formation  of  sulphuric  acid  by  the  oxidation  of  hydrogen  sulphide,  espec- 
ially by  means  of  ferric  chloride.  As  much  as  15  mg.  of  barium  may  be  pre- 
cipitated in  this  manner.  A  still  further  loss  of  15  mg.  of  barium  as  well  as 
considerable  quantities  of  calcium  and  strontium  as  carbonates  may  occur 
during  the  precipitation  of  the  iron  group.  Smaller  quantities  may  be  pre- 
cipitated as  sulphate  or  carbonate  with  the  ammonium  sulphide  group.  (Curt- 
man  and  Frankel,  /.  Am.  Soc.,  33,  724  (1911).) 

If  large  quantities  of  the  metals  in  the  preceding  groups;  especially  iron, 
are  present,  barium,  strontium  and  calcium  may  fail  to  be  detected  for  this 
reason,  by  the  ammonium  carbonate  method  of  separation.  The  method  of 
Curtman  and  Frankel  should  then  be  used.  See  §197. 

2.  Do  not  boil  after  the  addition  of  ammonium  carbonate,  as  this  will  drive  off 


§194,  .;.  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES. 

ammonium  hydroxide  and  carbonate,  increasing  the  solubility  of  the  CaCO3  (note  3 
and  §178). 

3.  The  precipitation  of  bivium,  strontium  and  calcium  by  ammonium  car- 
bonate in  the  presence  of  ammonium  chloride  is  not  as  complete  as  would  be  desir- 
able in  v/ry  delicate  analyses  The  carbonates  of  barium,  strontium  and  calcium 
are  all  slightly  soluble  in  ammonium  chloride  solution;  and  while  the  prescribed 
addition  of  ammonium  hydroxide,  and  excess  of  ammonium  carbonate,  greatly  re- 
duces the  solubility  of  the  precipitat  d  carbonates,  yet  even  with  th  se  the  precipi- 
tation is  not  absolute,  though  more  nearly  so  with  strontium  than  with  barium  and 
c  tlcium.  i  hus,  in  c/umtit  ti»e  analyses,  if  barium  and  calcium  are  precipitated  as 
<  urbonates,  it  must  be  done  in  the  absence  of  ammonium  chloride  or  sulphate,  and 
the  precipitate  washed  with  water  containing  ammonium  hydroxide.  As  much  as 
10  mg.  of  barium  may  remain  in  solution  in  the  presence  of  much  ammonium 
salts.  (Curtman  and  Frankel,  J.  Am.  Soc.,  33,  724  (1911).) 

4-  If  barium  be  absent,  as  evidenced  by  the  failure  to  obtain  a  precipitate 
with  K2Cr20T  ,  the  solution  may  at  once  be  divided  into  two  portions  to  test  for 
strontium  and  calcium. 

5.  With  care  the  reprecipitation  by  ammonium  carbonate,  for  the   separa- 
tion from  the  excess  of  K,Cr,O7  ,  may  be  neglected  and  the  filtrate  from  the 
barium,  yellow,  at  once  divided  into  two  portions  and  tested  for  Sr  and  Ca  . 
lleprecipitation  always  causes  the  loss  of  some  of  the  metals,  due  to  the  solu- 
bility of  the  carbonates  in  the  ammonium  acetate  formed.     On  the  other  hand, 
traces  may  escape  observation  in  the  yellow  chromate  solution. 

6.  Before  reprecipitation  with  (NH4)iCO3  ,  an  excess  of  ammonium  hydroxide 
should  be  added  to  prevent  the   liberation  of  CO2   when  the  ammonium  car- 
bonate is  added. 

7.  Strontium  sulphate  is  so  sparingly  soluble  in  water    (§187,   5c)   that  its 
precipitation  by  CaSO4    (or  other  sulphates  in  absence  of  Ca)    is   sufficiently 
delicate  to  detect  very  small  amounts  of  that  metal.     However,  it  is  sufficiently 
soluble  in  water  to  serve  as  a  valuable  reagent  to  detect  the  presence  of  traces 
of  barium.     Obviously   SrS04   will   not   precipitate   solutions   of   calcium    salts. 
Solutions  of  strontium  and  barium  salts    (except  SrSO4)    are  all  precipitated 
by  CaSO4.     The  presence  of  excess  of  calcium  salts  lessens  the  delicacy  of  the 
precipitation  of  strontium  salts  by  calcium  sulphate. 

8. — In  very  dilute  solutions  the  sulphates  of  the  alkaline  earths  are  not 
precipitated  rapidly.  Time  should  be  allowed  for  the  complete  precipitation. 
Boiling  and  evaporation  facilitates  the  reaction. 

9.  It  should  be  noticed  that  the  test  for  calcium  as  an  oxalate  is  made  upon 
that  portion  of  the  calcium  not  removed  by  (NH4)2SO4;  or  in  other  words  upon 
a  solution  of  CaSO,  (1-287).  A  solution  of  SrSO4  (1-10,000)  may  be  present, 
but  is  not  precipitated  by  i  NHi)2C2O4  .  The  presence  of  a  great  excess  of  (NH,)2SO4 
prevents  the  precipitation  of  traces  of  calcium  salts  by  (NH4)2C2O4 . 

§193.  Manipulation. — To  a  portion  of  the  filtrate  from  the  carbonates 
of  Ba ,  Sr ,  arid  Ca  add  a  drop  or  two  of  (NH4)2S04.  A  slight  precipitate 
indicates  a  trace  of  barium.  To  the  filtrate  a  few  drops  of  (NH4)2C204 
are  added.  A  trace  of  calcium  is  indicated  by  a  slight  precipitate.  Filter 
if  a  precipitate  is  obtained  and  test  the  filtrate  for  Mg  with  Na2HP04 . 
A  white  crystalline  precipitate — MgNH4P04 — is  evidence  of  the  presence 
of  magnesium.  The  other  portion  of  the  filtrate  from  the  carbonates  of 
Ba ,  Sr ,  and  Ca  is  reserved  to  be  tested  for  the  alkali  metals  (§211). 

§194 .  No'es. — 1 .  B  y  some,  magnesium  is  classed  in  the  last  or  alkali  group  instead 
of  in  the  alkaline  earth  group.  It  is  not  precipitated  by  the  (NH^COs,  yet  in  the 
general  properties  of  its  salts  it  is  so  closely  related  to  Ba,  Sr  and  Ca,  that  it  is 
much  better  regarded  as  a  subdivisi  n  of  that  group  than  as  belonging  to  the  alkali 
group  ( §  7  >  and  ff. ) . 

2.  Traces  of  t?a ,  Sr  and  Ca  may  remain  in  solution  after  adding  (NH4)2CO3 
and  warming;  due  to  the  solvent  action  of  the  ammonium  salts  present.  To  pre- 


226 


SEPARATION  OF  BARIUM,  STRONTIUM,   AND  CALCIUM.    §194,  3. 


vent  the.se  traces  giving  a  test  lor  magnesium  with  Na2HPO4  a  drop  or  two  of 
(NH4)2SO4  is  added  to  remove  barium  or  strontium  and  a  few  drops  of  (NH4)2C2O4 
to  remoye  calcium.  The  precipitate  (if  any  forms)  is  removed  by  nitration,  before 
the  Na2HPO4  is  added. 

3.  The  precipitate  of  P^g  '•>  V  "4  does  not  always  form  rapidly,  if  only  small 
amounts  of  Mg  are  present,  and  the  solution  should  be  allowed  to  stand.    Rubbing 
the  sides  of  the  test  tube  with  a  glass  stirring  rod  promotes  the  precipitation. 

4.  The  precipitation  of  Mg  as  MgtftI4FO4  is  fairly  delicate  (1-71,492)  (Kissel, 
Z.,  1F69,  8,  173  ;   but   not   very  characteristic,  as  the  phosphates  of  nearly  all  the 
metals  are  white  and  insoluble  in  water.     Hence  the  reliability  of  this  test  for  mag- 
nesium depends  upon  the  rigid  exclusion  of  the  other  metals  (not  alkalis)  by  the  pre- 
vious processes  of  analysis.     The  precipitate  should  be  carefully  examined  with  a 
magnifying  glass  to  ascertain  if  it  is  crystalline  which  is  characteristic  of  magnesium 
ammonium  phosphate.     Small  amounts  of  aluminum  may  be  present  at  this  point 
and  a  small  amorphous  precipitate  may  be  aluminum  phosphate.     It  should  be 
filtered  off  and  treated  with  a  few  drops  of  acetic  acid  which  readily  dissolves  mag- 
nesium ammonium  phosphate,  but  not  aluminum  phosphate.     The  filtrate  should 
be  neutralized  with  ammonia  and  a  few  drops  of  Na2HPO4  added.     On  standing 
for  some  time  small  transparent  crystals  are  deposited  on  the  walls  of  the  test  tube 
if  magnesium  is  present. 

5.  Lithium  phosphate  is  not  readily  soluble  in  water  or  ammonium  salts  and 
may  give  a  test  for  magnesium.     See  §210,  Gd. 

§195.  The  unlike  solubilities  in  alcohol,  of  the  chlorides  and  nitrates  of 
barium,  strontium  and  calcium  enable  us  to  separate  them  very  closely  by 
absolute  alcohol,  and  approximately  by  "  strong  alcohol,"  as  follows: 

Dissolve  the  carbonate  precipitate  in  HC1 ,  evaporate  to  dry  ness  on  the 
water-bath,  rub  the  residue  to  a  fine  powder  in  the  evaporating  dish,  and 
digest  it  with  alcohol.  Filter  through  a  small  filter,  and  wash  Avith  alcohol 
(56-,  §§186,  187  ami  188). 


Residue:  BaCl2. 

Dissolve  in  water 
test  with  CaSO4, 
SrSO4  ,  K2Cr2O7, 
etc. 

Filtrate:  SrCl2  and  CaCl2  . 

Evaporate  to  dryness,  change  to  nitrates  by  adding  a  few  drops 
of  HNO3  .     Evaporate  the  nitrates  to  dryness,  powder,  digest 
with  alcohol,*  filter  and  wash  with  alcohol  (or  digest  and  wash 
with  equal  volumes  of  alcohol  and  ether.) 

Residue:   Sr(NO3)2. 
Precipitation  by  CaSO4  in  water 
solution:    flame  test,  etc. 

Filtrate:    Ca(NO3)2. 
Precipitation  by  H2SO4  in  alco- 
hol solution,  by  (NH4)2C2O4, 
etc. 

Or,  the  alcoholic  filtrate  of  SrCl2  and  CaCl2  may  be  precipitated  with  (a 
drop  of)  sulphuric  acid,  the  precipitate  filtered  out  and  digested  with 
solution  of  (NH4),S04  and  a  little  NH4OH .  Residue,  SrS04 .  Solution 
contains  CaS04 ,  precipitable  by  oxalates. 

§196.  If  the  alkaline  earth  metals  are  present  in  the  original  material 
as  phosphates,  or  in  mixtures  such  that  the  treatment  for  solution  will 
bring  them  in  contact  with  phosphoric  acid;  the  process  of  analysis-  must 
be  modified.  One  of  the  methods  given  under  analysis  of  third  and  fourth 
group  metals  in  presence  of  phosphate*  (§145  and//.)  must  be  employed. 

§197.  The  presence  of  oxalates  will  also  interfere,  necessitating  the 
evaporation  and  ignition  to  decompose  the  oxalic  acid  (§151). 


*  Instead  of  alcohol  the  residue  of  the  nitrates  may  be  boiled  with  amyl  alcohol.    Calcium 
nitrate  is  dissolved  making  a  complete  separation  from  the  strontium  nitrate  '§188,  5c), 


§200.  THE    ALKALI   GROUP.  227 

THE  ALKALI  GROUP  (SIXTH  GROUP). 
Potassium.    K  =  ;J9.10.     Caesium.     Cs  =  132.81. 
Sodium.    Na  =  23.00.     Rubidium.     Rb  =  85.45. 
Ammonium.     (NH4)'.     Lithium.    Li  =  6.94. 

§198.  The  metals  of  ilio  alkalis  are  highly  combustible,  oxidizing  qnickly 
in  the  air,  displacing  the  hydrogen  of  water  even  more  rapidly  than  zinc 
or  iron  displaces  the  hydrogen  of  acids,  and  displacing  non-alkali  metals 
from  their  oxides  and  salts.  As  elements  they  are  very  strong  reducing 
agents,  while  their  compounds  are  very  stable,  and  "not  liable  to  either  re- 
duction or  oxidation  by  ordinary  means.  The  five  metals,  Cs ,  Rb ,  K , 
Na ,  Li ,  present  a  gradation  of  electro-positive  or  basic  power,  cesium 
being  strongest,  and  the  others  decreasing  in  the  order  of  their  atomic 
weights,  lithium  decomposing  water  with  less  violence  than  the  others. 
Their  specific  gravities  decrease,*  their  fusing  points  rise,  and  as  carbon- 
ates their  solubilities  lessen,  in  the  same  order.  In  solubility  of  the  phos- 
phate, also,  lithium  approaches  the  character  of  an  alkaline  earth  (§6). 

Ammonium  is  the  basal  radical  of  ammonium  salts,  and  as  such  has 
many  of  the  characteristics  of  an  alkali  metal.  The  water  solution  of  the 
gas  ammonia,  NH3  (an  anhydride),  from  analogy  is  supposed  to  contain 
ammonium  hydroxide,  NH4OH,  known  as  the  volatile  alkali.  Potassium 
and  sodium  hydroxides  are  the  fixed  alkalis  in  common  use. 

§199.  The  alkalis  are  very  soluble  in  water,  and  all  the  important  salts 
of  the  alkali  metals  (including  NH4)  are  soluble  in  water,  not  excepting  their 
carbonates,  phosphates  (except  lithium),  and  silicates;  while  all  other 
metals  form  hydroxides  or  oxides,  either  insoluble  or  sparingly  soluble,  and 
carbonates,  phosphates,  silicates,  and  certain  other  salts  quite  insoluble  in 
water. 

Their  compounds  being  nearly  all  soluble,  the  alkali  metals  are  not  pre- 
cipitated by  ordinary  reagents,  and,  with  few  exceptions,  their  salts  do  not 
precipitate  each  other.  In  analysis,  they  are  mostly  separated  from  other 
metals  by  non-precipitation. 

§200.  In  accordance  with  the  insolubility  in  water  of  the  non-alkali 
hydroxides  and  oxides,  the  alkali  hydroxides  precipitate  all  non-alkali  metals, 
except  that  ammonium  hydroxide  does  not  precipitate  barium,  strontium, 
and  calcium.  These  precipitates  are  hydroxides,  except  those  of  mercury, 
silver,  and  antimony.  But  certain  of  the  non-alkali  hydroxides  and 
oxides,  though  insoluble  in  water,  dissolve  in  solutions  of  alkalis;  hence, 
when  added  in  excess,  the  alkalis  redissolve  the  precipitates  they  at  first  pro- 
duce with  salts  of  certain  metals,  viz. :  the  hydroxides  of  Pb ,  Sn ,  Sb  (oxide), 

*  Except  those  of  potassium  (0,875)  and  sodium  (0.9735). 


228  POTASSIUM.  §201. 

Zn ,  Al,  and  Cr  dissolve  in  the  fixed  alkalis;  and  oxide  of  Ag  and  hy- 
droxides of  Cu ,  Cd ,  Zn ,  Co ,  and  Ni  dissolve  in  the  volatile  alkali. 

§201.  Solutions  of  the  alkalis  are  caustic  to  the  taste  and  touch,  and 
turn  red  litmus  blue;  also,  the  carbonates,  acid  carbonates,  normal  and 
dibasic  phosphates,  and  some  other  salts  of  the  alkali  metals,  give  the 
"  alkaline  reaction "  with  test  papers.  Sodium  nitrof  erricyanide,  with 
hydrogen  sulphide,  gives  a  delicate  reaction  for  the  alkali  hydroxides 
(§207,  66). 

§202.  The  hydroxides  and  normal  carbonates  of  the  alkali  metals  are  not 
decomposed  by  heat  alone  (as  are  those  of  other  metals),  and  these  metals 
form  the  only  acid  carbonates  obtained  in  the  solid  state. 

§203.  The  fixed  alkalis,  likewise  many  of  their  salts,  melt  on  platinum 
foil  in  the  flame,  and  slowly  vaporize  at  a  bright  red  heat.  All  salts  of 
ammonium,  by  a  careful  evaporation  of  their  solutions  on  platinum  foil, 
may  be  obtained  in  a  solid  residue,  which  rapidly  vaporizes,  wholly  or 
partly,  below  a  red  heat  (distinction  from  fixed  alkali  metals). 

§204.  The  hydroxides  of  the  fixed  alkali  metals,  and  those  of  their  salts 
most  volatile  at  a  red  heat,  preferably  their  chlorides,  impart  strongly 
characteristic  colors  to  a  non-luminous  flame,  and  give  well-defined  spectra 
with  the  spectroscope. 

§205.  Potassium.    K  =  39.10.    Valence  one. 

1.  Properties— Specific  gravity,   0.875  at  13°  (Baumhauer,  B.,  1873,  6,  655). 
Melting  point,   62.3°   (Cir.   B.  of  S.,   1915).     Boiling  point,   719°  to  731°   (Car- 
nelley  and  Williams,  B.,   1879,   12,   1360);  667°    (Perman,  J.   C.,   1889,   55,   328). 
Silver-white  metal  with  a  bluish  tinge.     At  ordinary  temperature  of  a.  wax-like 
consistency,  ductile  and  malleable;  at  0°  it  is  brittle.     It  is  harder  than  Na 
and  is  scratched  by  Li  ,  Pb  ,  Ca  and  Sr  .     The  glowing  vapor  is  a  very  beautiful 
intense  violet  (Dudley,  Am.,  1892,  14,  185).     It  is  next  to  caesium  and  rubidium, 
the  most  electro-positive  of  all  metals,  remains  unchanged  in  dry  air,  oxidizes 
rapidly   in    moist    air,    and    decomposes    water    with    great    violence,    evolving 
hydrogen,    burning   with    a    violet    flame.       At    a    red    heat    CO    and    C02    are 
decomposed,  at  a  white  heat  the  reverse  action  takes  place.     Liquid  chlorine 
does  not  attack  dry  potassium  (Gautier  and  Charpy,  C.  r.,  1891,  113,  597).     Acids 
attack  it  violently,  evolving  hydrogen. 

2.  Occurrence. — Very  widely  distributed  as  a  portion  of  many  silicates.     In 
sea  water  in  small  amount  as  KC1 .     In  numerous  combinations  in  the  large 
salt   deposits,    especially    at    Stassfurt;    e.g.,    carnallite,    KCl.Mg-Cl,    +    6H2O; 
kainite,  K2S04.MgSO4.MgCl2    +   6H20  ,   etc.     As  an   important  constituent  of 
many  plants — grape,  potato,  sugar-beet,  tobacco,  fumaria,  rumex,  oxalis,  etc. 

3.  Preparation. — (1)    By   reduction   of   the   carbonate   with   carbon.     (2)    By 
electrolysis    of   the   hydroxide    (Horning   and    Kasemeyer,    B.,    1889,    22,   277c; 
Castner,  B.,  1892,  25,  179c).     (3)  By  reduction  of  K2CO3  or  KOH  with  iron  car- 
bide: 6KOH  +  2FeC2  =  6K  +  2Fe  +  2CO  +  2C02  +  3H,   (Castner,  C.  N.,  1886, 
54,   218).     (//)    By  reduction   of  the   carbonate   or   hydroxide  with   Fe   or   Mg 
(Winkler,  B.,  1890,  23,  44). 

4.  Oxides  and  Hydroxide. — Potassium  oande*  K20  ,  is  prepared  by  carefully 

*  The  existence  of  tfce  oxides  M',Q  of  K,  Na  and  Jito  is  disputed  (Erdmann  apd  Koetlmer,  4* 
1896,  294,  55), 


§205, 65.  POTASSIUM.  229 

heating1  potassium  with  the  necessary  amount  of  oxygen  (air)  (Kuhnemann, 
C.  C.,  1863,  491);  also  by  heating-  K,O4  with  a  mixture  of  K  and  Ag  (Beketoff, 
C.  C.,  1881,  643).  It  is  a  hard,  gray  mass,  melting  above  a  red  heat.  \Yater 
changes  it  to  KOH  with  generation  of  much  heat.  I'oldxxiuin  lii/dro.ridc,  KOH, 
is  formed  by  treating  K  or  K2O  with  water;  by  boiling  a  solution  of  K,CO:! 
with  Ba  ,  Sr  or  Ca  oxides;  by  heating  K,C03  with  Fe.Og  to  a  red  heat  and 
decomposing  the  potassium  ferrate  with  water  (Ellershausen,  C.  C.,  1891,  (1), 
1047;  (2),  399).  Pure  water-free  KOH  is  a  white,  hard,  brittle  mass,  melting 
at  a  red  heat.  It  dissolves  in  water  with  generation  of  much  heat.  Potassium 
supcro.t'idc,  K._>O4  ,  is  formed  when  K  is  heated  in  contact  with  abundance  of  air 
(Harcoiirt.  ,/.  (7.,  1862,  14,  267);  also  by  bringing  K  in  contact  with  KNO3 
heated  until  it  begins  to  evolve  0  (liolton,  C.  A'.,  1886,  53,  289).  It  is  an  amor- 
phous powder  of  the  color  of  lead  chromate.  Upon  ignition  in  a  silver  dish 
oxygen  is  evolved  and  K2O  and  AgoO  formed  (Harcourt,  I.  c.).  Moist  air  or 
water  decomposes  it  with  evolution  of  oxygen.  It  is  a  powerful  oxidizing 
agent,  oxidizing  S°  to  Svi  ,  P°  to  PV  ,  K  ,  As  ,  Sb  ,  Sn  ,  Zn  ,  Cu  ,  Fe  ,  Ag  and  Pt 
to  the  oxides  (Bolton,  I.  c.'  Brodie,  Proc.  Roy.  &oc.,  1863,  12,  209). 

5.  Solubilities. — K  and  K20  dissolve  in  water  with  violent  action,  forming 
KOH  ,  which  reacts  with  all  acids  forming  soluble  salts.  Potassium  dissolves 
in  alcohol,  forming  potassium  alcoholate  and  hydrogen. 

Potassium  platinum  chloride,  acid  tartrate,  silico-fluoride,  picrate,  phos- 
phomolybdate,  perchlorate,  and  chlorate  are  only  sparingly  soluble  in 
cold  water,  and  nearly  insoluble  in  alcohol.  The  carbonate  and  sulphate 
are  insoluble  in  alcohol.  , 

6.  Reactions,  a. — Potassium  and  sodium  hydroxides  are  very  strong 
bases,  fixed  alkalis,  and  precipitate  solutions  of  the  salts  of  all  the  other 
metals  (except  Cs ,  Rb ,  and  Li),  as  oxides  or  hyHroxides.  These  precipi- 
tates are  quite  insoluble  in  water,  except  the  hydroxides  of  Ba ,  Sr ,  and 
Ca .  Excess  of  the  reagent  causes  a  resolution  with  the  precipitates  of 
Pb,  Sb,  Sn,  Al,  Cr,  and  Zn ,  forming  double  oxides  as,  K2Pb02 ,  potas- 
sium plumbite,  etc.  Potassium  carbonate  is  deliquescent,  strongly  alkaline, 
and  precipitates  solutions  of  the  salts  of  the  metals  (except  Cs ,  Rb ,  Na  , 
and  Li),  forming  normal  carbonates  with  Ag ,  Hg',  Cd ,  Fe",  Mn ,  Ba ,  Sr , 
and  Ca  ;  oxide  with  Sb  ;  hydroxide  with  Sn ,  Fe"',  Al ,  Cr'"  and  Co'";  basic 
salt  with  Hg",  and  a  basic  carbonate  with  the  other  metals. 

&.--The  potassium  salts  of  HCN,  H4Fe(CN)6 ,  H3Fe(CN)6 ,  and  HCNS 
find  extended  application  in  the  detection  and  estimation  of  many  of  the 
heavy  metals. 

Tartaric  acid,  H2C4H406 ,  or  more  readily  sodium  hydrogen  tartrate, 
NaHC4H406 ,  precipitates,  from  solutions  sufficiently  concentrated,  potas- 
sium hydrogen  tartrate,  KHC4H400 ,  granular-crystalline.  If  the  solution 
be  alkaline,  acetic  or  tartaric  acid  should  be  added  to  strong  acid  reaction. 
The  test  must  be  made  in  absence  of  non-alkali  bases.  The  precipitate  is 
increased  by  agitation,  and  by  addition  of  alcohol.  It  is  dissolved  by 
fifteen  parts  of  boiling  water  or  eighty-nine  parts  water  at  25°,  by  mineral 
acids,  by  solution  of  borax,  and  by  alkalis,  which  form  the  more  soluble 
normal  tartrate,  K.,C4H406 ,  but  not  by  acetic  acid,  or  at  all  by  alcohol 
of  fifty  per  cent. 


230  POTASSIUM.  §205,  60. 

Picric  acid,  C6H2(N02)3OH ,  or  preferably  its  sodium  salt,  precipitates, 
from  solutions  not  very  dilute,  the  yellow,  crystalline  potassium  picrate, 
C6H2(N02)3OK,  soluble  in  260  parts  of  water  at  15°  C.  (Reichard, 
Z.  40,  25),  insoluble  in  alcohol,  by  help  of  which  it  is  formed  in  dilute 
solutions.  The  solution  must  be  nearly  neutral  to  avoid  precipitation  of 
the  slightly  soluble  picric  acid  (soluble  in  160  parts  water).  The  dried 
precipitate  detonates  strongly  when  heated. 

c. — If  a  neutral  solution  of  a  potassium  salt  be  added  to  a  solution  of  cobaltic 
nitrite,*  a  precipitate  of  the  double  salt  potassium  cobaltic  nitrite,  K3Co(NO2)6  , 
will  be  formed.  In  concentrated  solutions  the  precipitate  forms  immediately, 
dilute  solutions  should  be  allowed  to  stand  for  some  time;  sparingly  soluble  in 
water,  insoluble  in  alcohol  and  in  a  solution  of  potassium  salts,  hence  the 
precipitation  is  more  valuable  as  a  separation  of  cobalt  from  nickel  than  as  a 
test  for  potassium  (§132,  6c). 

Potassium  nitrate  is  not  found  abundantly  in  nature,  but  is  formed  by  the 
decomposition  of  nitrogenous  organic  substances  in  contact  with  potassium 
salts,  "  saltpeter  plantations  ";  or  by  treating  a  hot  solution  of  NaN03  with 
KC1  (Z).,  2,  2,  72).  It  finds  extended  application  in  the  manufacture  of  gun- 
powder, d. — See  §206,  6d. 

e. — Potassium  sulphide  may  be  taken  as  a  type  of  the  soluble  sulphides 
which  precipitate  solutions  of  the  metals  of  the  first  four  groups  as 
sulphides  except:  Hg'  becomes  HgS  and  Hg°,  Fe'"  becomes  FeS  and  S, 
and  Al  and  Cr  form  hydroxides.  The  sulphides  of  arsenic,  antimony  and 
tin  dissolve  in  an  excess  of  the  reagent,  more  rapidly  if  the  alkali  sulphide 
contain  an  excess  of  sulphur.  For  the  general  action  of  H,S  or  soluble 
sulphides  as  a  reducing  agent  see  the  respective  metals.  Potassium  sul- 
phate  is  used  to  precipitate  barium,  strontium,  and  lead.  It  almost  always 
occurs  in  nature  as  double  salt  with  magnesium,  K2S04.MgS04.MgCl2  -{- 
6H20 ,  kainite,  and  is  used  in  the  manufacture  of  KA1(S04)2 ,  K2C03  and 
KOH  .  As  a  type  of  a  soluble  sulphate  it  precipitates  solutions  of  lead, 
mercurosum,  barium,  strontium,  and  calcium;  calcium  and  mercurosum 
incompletely. 

/. — Potassium  chloride  precipitates  the  metals  of  the  first  group,  acting 
thus  as  a  type  of  the  soluble  chlorides.  It  is  much  used  with  sodium 
nitrate  in  the  preparation  of  potassium  nitrate  for  the  manufacture  of 
gunpowder,  in  the  preparation  of  K2C03 ,  KOH ,  and  also  as  a  fertilizer. 
Potassium  bromide  as  a  type  of  the  soluble  bromides  precipitates  solutions 
of  Pb,  Ag,  and  Hg  (Hg"  incompletely).  Potassium  iodide  finds  extended 
use  in  analytical  chemistry  in  that  it  forms  many  soluble  double  iodides; 
it  is  also  extensively  used  in  medicine.  As  a  typo  of  a  soluble  iodide  it 
precipitates  solutions  of  the  salts  of  Pb ,  Ag,  Hg  ,  and  Cu'.  Cu"  salts 
are  precipitated  as  Cul  with  liberation  of  iodine.  Fe"'  salts  are  merely 

*  One  cc.  of  cobaltous  nitrate  solution  and  three  cc.  of  acetic  acid  are  added  to  five  cc.  of  a  ten 
per  cent  solution  of  sodium  nitrite.  This  gives  a  yellowish  solution  having  an  odor  of  nitrous 
acid. 


§205,  7.  POTASSIUM.  231 

reduced   to  Fe"  salts  with  liberation  of  iodine.      Arsenic  acid  is  merely 

reduced  to  arsenous  acid  with  liberation  of  iodine. 

Potassium  chlorate  is  used  as  a  source  of  oxygen  and  as  an  oxidizing  agent 
in  acid  solutions.  Sodium  perchlorate,  NaClO4  ,  precipitates  from  solutions  of 
potassium  salts  i>ot<ixxinin  ftcn-hlonite,  KC1O.,  ,  sparingly  soluble  in  water  and 
almost  insoluble  in  strong  alcohol  (Kreider,  %.  anory.,  1895,  9,  342).  Potassium 
iodate  is  used  as  a  reagent  in  the  detection  of  barium  as  Ba(IO8)o  .  <j.  —  The 
oxides  of  iirxcuic  act  as  acid  anhydrides  toward  KOH  and  form  stable  soluble 
potassium  salts,  arsenites  and  arsenates,  which  react  with  the  salts  of  nearly 
all  the  heavy  metals,  h.  —  Potassium  chromate  and  dichromate  are  both  exten- 
sively used  as  "reagents,  especially  in  the  analysis  of  Ag  ,  Pb  and  Ba  salts. 

i.  —  Fluosilicic  acid,,  H2SiF6  ,  precipitates  from  a  neutral  or  slightly 
acid  solution  of  potassium  salts,  potassium  fluosilicate  (silico-fluoride), 
K,SiFc,  soluble  in  833.1  parts  of  water  at  17.5°;  in  104.8  parts  at  100°; 
and  in  327  parts  of  9.6  per  cent  HC1  at  14°  (Stolba,  J.  pr.,  1868,  103,  396\ 
The  precipitate  is  white,  very  nearly  transparent. 

;'.  —  Hydrochlorplatinic  acid,  H2PtClG  ,  added  to  neutral  or  acid  solutions 
not  too  dilute,  precipitates  potassium  clilorplatinate,  K2PtCl0  ,  crystalline, 
yellow.  Non-alkali  bases  also  precipitate  this  reagent,  and  if  present  must  be 
removed  before  this  test.  The  precipitate  is  soluble  in  19  parts  of  boiling 
water,  or  111  parts  of  water  at  10°.  Minute  proportions  are  detected  by 
evaporating  the  solution  with  the  reagent  nearly  to  dryness,  on  the  water- 
bath,  and  then  dissolving  in  alcohol;  the  yellow  crystalline  precipitate, 
octahedral,  remains  undissolved,  and  may  be  identified  under  the  micro- 
scope. 


k.  —  An  alcoholic  solution  of  BiCl3  in  excess  of  Na^Os  gives  a  yellow  precipitate 
with  solutions  of  potassium  salts  (Pauly,  C.  C.,  1887,  553).  1.  —  Gold  chloride  added 
to  sodium  and  potassium  chloride  forms  chloraurates,  e.g.,  KAuCl,  -f  !'H2O  .'  If 
these  salts  are  dried  at  100°  to  110°  to  remove  water  and  acids,  the  sodium  salt  is 
soluble  in  ether  (separation  from  potassium)  (Fasbender,  C.  C.,  1894,  1,  409). 

7.  Ignition.  —  Ignited  potassium  hydroxide  or  potassium  carbonate  is  a 
valuable  desiccating  agent  for  use  in  desiccators  or  in  liquids.  A  mixture 
of  molecular  proportions  of  K2COa  and  Na2C03  melts  at  a  lower  tempera- 
ture than  either  of  the  constituents,  and  is  frequently  employed  in  fusion 
for  the  transposition  of  insoluble  metallic  compounds  :  BaS04  -f-  K>C03  = 
BaC03  -f  K2S04  . 

Potassium  compounds  color  the  flame  violet.  A  little  of  the  solid 
substance,  or  residue  by  evaporation,  moistened  with  hydrochloric  acid, 
is  brought  on  a  platinum  wire  into  a  non-luminous  flame.  The  wire 
should  be  previously  washed  with  HC1,  and  held  in  the  flame  to  insure 
the  absence  of  potassium.  The  presence  of  very  small  quantities  of 
sodium  enables  its  yellow  flame  completely  to  obscure  the  violet  of  potas- 
sium; but  owing  to  the  greater  volatility  of  the  latter  metal,  flashes  of 
violet  are  sometimes  seen  on  the  first  introduction  of  the  wire,  or  at  the 
border  of  the  flame,  or  in  its  base,  even  when  enough  sodium  is  present 
to  conceal  the  violet  at  full  heat.  The  interposition  of  a  blue  glass,  or 


232  SODIUM.  §205,  8. 

to  conceal  the  violet  at  full  heat.  The  interposition  of  a  blue  glass,  or 
prism  filled  with  indigo  solution,  sufficiently  thick,  entirely  cuts  off  the 
yellow  light  of  sodium,  and  enables  the  potassium  name  to  be  seen.  The 
red  rays  of  the  lithium  flame  are  also  intercepted  by  the  blue  glass  or 
indigo  prism,,  a  thicker  stratum  being  required  than  for  sodium.  It 
organic  substances  are  present,  giving  luminosity  to  the  flame,,  they  must 
be  removed  by  ignition.  Certain  non-alkali  bases  interfere  with  the 
examination.  Silicates  may  be  fused  with  pure  gypsum,  giving  vapor  of 
potassium  sulphate.  Bloxam  (J.  C.,  1865,  18,  "229)  recommends  to  fuse 
insoluble  alkali  compounds  with  a  mixture  of  sulphur,  one  part,  and 
barium  nitrate,  six  parts;  cool,  dissolve  in  w^ater,  remove  the  barium  with 
NH4OH  and  (NH4)2C03  and  test  for  the  alkalis  as  usual. 

The  volatile  potassium  compounds,  when  placed  in  the  flame,  give  a 
widely-extended  continuous  spectrum,  containing  two  characteristic  lines; 
one  line,  K  «,  situated  in  the  outermost  red,  and  a  second  line,  K  /?,  far  in 
the  violet  rays  at  the  other  end  of  the  spectrum. 

8.  Detection. — Potassium  is  usually  identified  by  the  violet  blue  color 
which  most  of  its  salts  impart  to  the  Bunsen  flame  (7).     Sodium  inter- 
feres but  the  intervention  of  a  cobalt  glass  (§132,  7)   or  a  solution  of 
indigo  cuts  out  the  yellow  color  of  the  sodium  flame  and  allows  the  violet 
of  the  potassium  to  be  seen.     Some  of  the  heavy  metals  interfere,  hence 
the  test  should  be  made  after  the  removal  of  the  heavy  metals  (§§211 
and  212). 

Potassium  may  be  precipitated  as  the  platinichloride  (6/);  as  the  per- 
chlorate  (6/);  as  the  silico-fluoride  (6i);  as  the  acid  tartrate  (66);  etc. 
Certain  of  these  reactions  are  much  used  for  the  quantitative  estimation 
(9)  of  potassium  but  are  seldom  used  for  its  detection  qualitatively. 

9.  Estimation. — (1)  Potassium  is  converted  into  the  sulphate  or  phosphate 
and  weighed  as  such.     (2)  It  is  precipitated  and  weighed  as  the  double  chloride 
with  platinum.     (3)  If  present  as  KOH  or  K2CO3  it  is  titrated  with  standard 
acid   (Kippenberger,  Z.  angew.,  1894,  495).     (4)   It  is  precipitated  with  H2SiF6 
and  strong  alcohol.     (5)   Indirectly  when  mixed  with  sodium,   by  converting 
into   the   chlorides   and  weighing   as   such;   then    determining   the   amount    of 
chlorine   and   calculating  the  relative  amounts  of   the   alkalis.     (6)    It  is  pre- 
cipitated  as   the   bitartrate  in   presence   of   alcohol    and,   after  nitration    and 
solution  in  hot  water,  titrated  with  deci-normal  KOH.     (7)  By  precipitation  as 
the  perchlorate,  KC1O4  (Wense,  Z.  angew.,  1892,  233;  Caspari,  Z.  angew.,  1893,  68). 

10.  Oxidation. — Potassium  is  a  very  powerful  reducing  agent,  its  affinity 
for  oxygen  at  temperatures  not  too  high  is  greater  than  that  of  any  other 
element  except  Cs  and  Rb  .  For  oxidizing  action  of  K204  see  4. 

§206.  Sodium.     Na  =  23.00.     Valence  one. 

1.  Properties.—  Specific  gravity,  0.9735  at  13.5°  (Baumhauer,  B.,  1873,  6, 
665);  0.7414  at  the  boiling  point  (Ramsay,  B.,  1880,  13,  2145).  Melting  point, 
97.5°  (Cir.  B.  of  S.,  1915).  Boiling  point,  742°  (Perman,  C.  N.,  1889,  59,  237) 


CALIFORNIA   COLLEQi 

§206, 6d.  SODIUM.  PHARMACY  233 

A  silver-white  metal  with  a  strong-  metallic  lustre.  At  ordinary  temperatures 
it  is  softer  than  Li  or  Pb,  and  can  be  pressed  together  between  the  fingers; 
at  — 20°  it  is  quite  hard;  at  0°  very  ductile.  It  oxidizes  rapidly  in  moist  air 
and  must  be  kept  under  benzol  or  kerosene.  It  decomposes  water  violently 
even  at  ordinary  temperatures,  evolving-  hydrogen,  which  frequently  ignites 
from  the  heat  of  the  reaction:  2Na  +  2H2O  =  2NaOH  +  H2  .  It  burns,  when 
heated  to  a  red  heat,  with  a  yellow  flame.  Pure  dry  Na  is  scarcely  at  all 
attacked  by  dry  HC1  (Cohen,  C.  N.,  1886,  54,  17). 

2.  Occurrence. — Never  occurs  free  in  nature,  but  in  its  various  combinations 
one  of  the  most  widely  diffused  metals.     There  is  no  mineral  known  in  which 
its   presence  has  not   been   detected.     It   occurs   in   all  waters   mostly   as   the 
chloride  from  traces  in  drinking  waters  to  a  nearly  saturated  solution  in  some 
mineral  waters  and  in  the  sea  water.     It  is  found  in  enormous  deposits  as  rock 
salt,  NaCI;  as  Chili  saltpeter,  NaNO3;  in  lesser  quantities  as  carbonate,  borate, 
sulphate,  etc. 

3.  Preparation. —  (1)   By  igniting-  the  carbonate  or  hydroxide  with  carbon; 
(2)   by  igniting-  the  hydroxide  with  metallic  iron;    (3)   by   electrolysis  of  the 
hydroxide;  (4)  by  gently  heating  the  carbonate  with  Mg- . 

4.  Oxides   and   Hydroxides. — Sodium    oxide,    Na.O  ,    is    formed    by    burning 
sodium  in  oxygen  or  in  air  and  heating  again  with  Na  to  decompose  the  Na2O2 
(§205,  4,  footnote).     Sodium  hydroxide,  NaOH  ,  is  formed  by  dissolving  the 
metal  or  the  oxide  in  water  (Rosenfeld,  J.  pr.,  1893,  (2),  48,  599);  by  treating 
a  solution  of  sodium  carbonate  with  lime;  by  fusion  of  NaN03  with  CaC03  , 
CaO  and  Na,CO3  are  formed  and  the  mass  is  then  exhausted  with  water:  by 
igniting  Na2C03  with  Fe2O3  ,   forming  sodium  ferrate,  which  is  then   decom- 
posed with  hot  water  into  NaOH  and  Fe(OH)3   (Solvay,  C.  C.,  1887,  829).     It  is 
a   white,    opaque,    brittle    crj-stalline    body,    melting    under    a    red    heat.     The 
fused  mass  has  a  sp.  gr.  of  2.13  (Filhol,  A.  Ch.,  1847,  (3),  21,  415).     It  has  a  very 
powerful  affinity  for  wrater,   gradually   absorbing  water  from   CaCl2    (Mtiller- 
Erzbach,  B.,  1878,  11,  409).     It  is  soluble  in  about  0.47  part  of  water  according 
to  Bineau  (C.  r.,  1855,  41,  509). 

Sodium  peroxide,  Na,02  ,  is  formed  by  heating  sodium  in  CO,,  free  air  or 
oxygen  (Prud'homme,  C.  C.,  1893,  (1),  199).  It  reacts  as  H202  ,  partly  reducing 
and  partly  oxidizing.  It  may  be  fused  without  decomposition.  Water  decom- 
poses it  partially  into  NaOH  and  H,O2  . 

5.  Solubilities. — Sodium  and  sodium  oxide  dissolve  in  water,  forming 
the  hydroxide,  the   former  with   evolution   of  hydrogen.     In  acids  the 
corresponding  sodium  salts  are  formed,  all  soluble  in  water  except  sodium 
pyroantimonate,  which  is  almost  insoluble  in  water,  and  the  fluosilicate 
sparingly  soluble. 

The  nitrate  and  chlorate  are  deliquescent.  The  carbonate  (10  aq),  sul- 
phate (10  aq),  sulphite  (8  aq),  phosphate  (12  aq),  and  the  acetate  (3  aq)  are 
efflorescent. 

6.  Keactions.     a. — As  reagents  sodium  hydroxide  and  carbonates  act  in 
all  respects  like  the  corresponding  potassium  compounds,  which  see. 

ft. — By  the  greater  solubility  of  the  picrate  and  acid  tnrlrafc  of  sodium,  that 
metal  is  separated  from  potassium  (§205,  6ft).  c. — Sodium  nitrate  occurs  in 
nature  in  large  quantities  as  Chili  saltpeter,  used  as  a  fertilizer,  for  the  manu- 
facture of  nitric  acid,  with  KC1  for  making  KNO3  ,  etc. 

d. — Sodium  phosphate,  Na.HP04 ,  is  much  used  as  a  reagent  in  the 
precipitation  and  estimation  of  Pb ,  Mn ,  Ba ,  Sr ,  Ca ,  and  Mg  .  The 
phosphates  of  all  metals  except  the  alkalis  are  insoluble  in  water  (lithium 
phosphate  is  only  sparingly  soluble  (§210,  5c),  soluble  in  acids).  Solu- 


234  SODIUM.  §206,  6e. 

tions  of  alkali  phosphates  precipitate  solutions  of  all  other  metallic  salts 
as  phosphates  (secondary,  tertiary  or  basic)  except:  HgCl2  precipitates  as 
a  basic  chloride  (§58,  6d),  and  antimony  as  oxide  or  oxychloride  (§70,  6d). 

e,  f,  g,  1i. — As  reagents  the  sodium  salts  react,  similar  to  the  corresponding- 
potassium  salts,  which  see.  i. — Sodium,  fluosilicate  is  soluble  in  153.3  parts 
H2O  at  17.5°  and  in  40.66  parts  at  100°  (Stolba,  £.,  1872,  11,  199);  hence  is  not 
precipitated  by  fluosilicic  acid  except  from  very  concentrated  solutions 
(separation  from  K).  j. — Sodium  chlorplatinate,  Na2PtCl6,  crystallizes  from  its 
concentrated  solutions  in  red  prisms,  or  prismatic  needles  (distinction  from  potassium 
or  ammonium).  A  drop  of  the  solution  to  be  tested  is  slightly  acidified  with  hydro- 
chloric acid  from  the  point  of  a  glass  rod  on  a  slip  of  glass,  treated  with  two  drops 
of  solution  of  chlorplatinic  acid,  left  a  short  time  for  spontaneous  evaporation  and 
crystallization,  and  observed  under  the  microscope,  k. — Sodium  picrate  soluble  in 
10  parts  of  water  is  used  as  a  reagent  for  potassium  salts  (Richard,  Z.,  40,  377). 

~k. — Solution  of  potassium  pyroantimonate,  K2H2Sb20T ,  produces  in 
neutral  or  alkaline  solutions  of  sodium  salts  a  slow-forming,  white,  crystal- 
line precipitate,  Na2H2Sb207 ,  almost  insoluble  in  cold  water.  The  reagent 
must  IDC  carefully  prepared  and  dissolved  when  required,  as  it  is  not  per- 
manent in  solution  (§70,  4c). 

7.  Ignition.— Sodium  bicarbonate,  NaHC03 ,  loses  H20  and  C02  at  125° 
becoming  Na2C03 ,  no  further  decomposition  till  400°  when  a  very  small 
amount  of  NaOH  is  formed  (Kirsling,  Z.  angew.,  1889,  332). 

Sodium  compounds  color  the  flame  intensely  yellow,  the  color  being 
scarcely  affected  by  potassium  (at  full  heat),  but  modified  to  orange-red 
by  much  lithium,  and  readily  intercepted  by  blue  glass.  Infusible  com- 
pounds may  be  ignited  with  calcium  sulphate.  The  test  is  interfered  with 
by  some  non-alkali  bases,  which  should  be  removed  (§§211  and  212). 

The  spectrum  of  sodium  consists  of  a  single  broad  band  at  the  D  line  in 
the  yellow  of  the  solar  spectrum  separable  into  two  bands,  Dy  and  D  „ ,  by 
prisms  of  higher  refractive  power. 

The  amount  of  sodium  in  the  atmosphere,  and  in  the  larger  number  of 
substances  designed  to  be  "  chemically  pure  "  is  sufficient  to  give  a  dis- 
tinct but  evanescent  yellow  color  to  the  flame  and  spectrum. 

8.  Detection. — Sodium  is  usually  detected  by  the  color  of  the  flame, 
yellow,  in  absence  of  the  heavy  metals.     In  the  usual  process  of  analysis 
the  presence   or  absence   of  sodium  is  determined  in  the   presence   of 
magnesium  (as  Na2HP04  is  the  usual  reagent  for  the  detection  of  mag- 
nesium, it  is  evident  that  the  presence  or  absence  of  the  sodium  must  be 
determined  before  the  addition  of  that  reagent);  and  as  that  metal  gives 
a  yellowish  color  to  the  flame  it  must  be  removed  if  small  quantities  of 
sodium  are  to  be  detected.     For  this  purpose  the  filtrate  from  Ba ,  Sr  and 
Ca  is  evaporated  to  dryness  and  gently  ignited  to  expel  all  ammonium 
salts;  then  taken  up  with  a  small  amount  of  water  and  the  magnesium 
precipitated  as  the  hydroxide  with  a  solution  of  barium  hydroxide.     After 


§207,   5.  AMMONIUM.  235 

filtration  the  barium  is  removed  by  (NH4)2C03  or  H2S04  and  the  filtrate 
tested  for  sodium  by  the  flame  or  by  the  pyroantimonate  test  (Glc). 

9.  Estimation. — (1)  If  present  as  hydroxide  or  carbonate,  by  titration  with 
standard  acid  (Lunge,  Z.  angew.,  1897,  41).  (2)  By  converting-  into  the  chloride 
or  sulphate  and  weighing  as  such.  (3)  In  presence  of  potassium  by  converting 
into  the  chloride,  weighing  as  such,  then  estimating  the  amount  of  chlorine 
with  AgN03  and  computing  the  amounts  of  K  and  Na .  (4  It  is  precipitated 
by  K,H2Sb2O7  and  dried  and  weighed  as  Na2H2Sbo07 . 

10.  Oxidation. — Sodium  ranks  with  potassium  as  a  very  powerful  re- 
ducing agent.  It  is  not  quite  so  violent  in  its  reaction  and  being  much 
cheaper  is  almost  universally  used  instead  of  potassium.  Sodium  peroxide 
may  act  both  as  a  reducing  and  oxidizing  agent.  The  action  is  similar  to 
H202  in  alkaline  solution,  which  see  (§244,  6). 


§207.  Ammonium.     (NH4)'.     Valence  one. 

1.  Properties. — Specific  gravity  of  NH3    gas,   0.589    (Fehling,  1,   384);  of  the 
liquid,  0.6234  at  0°    (Jolly,  A.,  1861,   117,  181).     The  liquid  boils  at  — 33.7°,  at 
0°  the  liquid  has  a  tension  of  4.8  atmospheres    (Bunsen,   Pogg.,   1839,  46,  95). 
Liquid  ammonia  is  a  colorless  mobile  liquid,  burns  in  air  when  heated  or  in 
oxygen  without  being  previously  heated.     At  ordinary  temperature  it  is  a  gas 
with  very  penetrating  odor.     It  burns  with  a  greenish-yellow  flame,  and  com- 
bines energetically  with  acids  to  form  salts,  the  radical  NH4  being  monovalent 
and  acting  in  many  respects  similar  to  K  and  Na  .     At  0°  one  volume  of  water 
absorbs  1049.6  volumes  of  the  gas;  at  15°,  727.22  volumes   (Carius,  A.,  1856,  99, 
144).     One    gram   of    water,    pressure    760   mm.    and    temperature    0°,    absorbs 
0.899  gram  of  NH3 ;  with  temperature  16°,  0.578  gram  (Sims,  A.,  1861,  118,  345). 

2.  Occurrence. — Free  ammonia  does  not  occur  in  nature.     Various  ammonium 
salts  occur  widely  distributed:  in  rain  water,  in  many  mineral  waters,  in  almost 
all  plants,  among  the  products  of  the  decay  or  decomposition  of  nitrogenous 
organic  bodies,  etc. 

3.  Preparation. — It  is  obtained  from  the  reduction  of  nitrates  or  nitrites  by 
nascent  hydrogen  in  alkaline  solution,  e.  g.,  8A1  +  5KOH  -f  3KNO3  -f  2H20  = 
SKAlOo  -f  3NH3;  by  the  reduction  with  the  hydrogen  of  the  zinc-copper  couple; 
by   boiling  organic   compounds   containing   nitrogen   with    KMn04    in    strong 
alkaline  solution   (as  in  water  analysis);  also  by  the  oxidation  of  nitrogen  in 
organic  bodies  with  strong  sulphuric  (Kjeldahl  method  of  nitrogen  determina- 
tion).    It  is  prepared  on  a  larger  scale  by  heating  an  ammonium  salt  with  lime 
(or  some  other  strong  base).     Nearly  all  the  ammonium  hydroxide  and  am- 
monium salts  of  commerce  are  obtained  as  a  by-product  in  the  production  of 
illuminating  gas  by  the  destructive  distillation  of  coal. 

4.  Hydroxide. — Ammonium   hydroxide,   NH4OH ,   is   made   by  passing 
ammonia,  NH3 ,  into  water.     The  gas  is  absorbed  by  the  water  with  great 
avidity,  and  a  strongly  alkaline  solution  is  produced.     A  solution  having 
a  sp.  gr.  of  0.90  at  15°  contains  28.33  per  cent  of  NH3  (Lunge  and  Wiernik, 
Z.  angeiu.,  1889,  183). 

5.  Solubilities.— Ammonia,  NH3 ,  and  all  ammonium  salts  are  soluble  in 
water.     Ammonia  dissolves  less  readily  in  a  strong  solution  of  potassium 
hydroxide  than  in  water.   The  carbonate  (acid),  and  phosphate  are  efflores- 
cent.    The  nitrate  and  acetate  are  deliquescent,  the  sulphate  slightly  deli- 
quescent. 


AMMONIUM.  §£07, 6a. 

6.  Reactions,  a. — The  fixed  alkali  hydroxides  and  carbonates  liberate 
ammonia,  NH3 ,  from  all  ammonium  salts,  in  the  cold  and  more  rapidly 
upon  heating.  Ammonium  hydroxide,  volatile  alkali,  colors  litmus  blue, 
neutralizes  acids,  forming  salts,  and  precipitates  solutions  of  the  metals  of 
the  first  four  groups,  manganese  and  magnesium  salts  imperfectly;  due  to 
the  solubility  of  the  hydroxide  formed,  in  the  ammonium  salt  produced 
by  the  reaction,  and  with  these  metals  if  excess  of  ammonium  salts  be 
present  no  precipitate  will  be  formed  by  the  NH4OH .  The  precipitate  is 
a  hydroxide  except:  with  Ag  and  Sb  it  is  an  oxide,  with  mercury  a  sub- 
stituted ammonium  salt  and  with  lead  a  basic  salt  (see  below,  fc  and  I). 
With  salts  of  Ag ,  Cu ,  Cd ,  Co ,  Ni ,  and  Zn  the  precipitate  redissolves  in 
excess  of  the  reagent.  Ammonium  carbonate,  (NN4)2C03 ,  is  unstable  and 
used  only  in  solution.  It  is  formed  by  adding  ammonium  hydroxide  to  a 
solution  of  the  acid  carbonate  of  commerce.  It  precipitates  solutions  of 
all  the  non-alkali  metals,  chiefly  as  carbonates  except  magnesium  salts 
which  are  not  at  all  precipitated,  as  a  soluble  double  salt  is  at  once  formed 
(separation  of  Ba ,  Sr ,  and  Ca  from  Mg).  With  salts  of  Ag ,  €u ,  Cd  ,  Co  , 
Ni ,  and  Zn ,  the  precipitate  is  redissolyed  by  an  excess  of  the  ammonium 
carbonate. 

&. — Dilute  solutions  of  picric  acid  with  ammonium  hydroxide  form  in- 
tensely colored  yellow  solutions,  a  precipitate  of  ammonium  picrate  being 
formed  if  the  solutions  are  quite  concentrated.  Tartaric  acid  precipitates 
ammonium  salts  very  closely  resembling  the  precipitate  of  potassium  acid 
tartrate.  The  ammonium  salt  is  more  soluble  in  water  than  the  potas- 
sium salt  and  does  not  leave  K2C03  upon  ignition.  Sodium  nitroferri- 
cyanide,  Na2Fe(NO)(CN)5 ,  added  to  a  mixture  of  NH4OH  and  H2S 
[(NH4)2S]  gives  a  very  intense  purple  color,  characteristic  of  alkali 
sulphides  and  the  manipulation  may  be  modified  so  as  to  give  a  very  deli- 
cate test  for  the  presence  of  an  alkali  hydroxide  or  of  hydrosulphuric  acid. 
In  no  case,  however,  can  the  H2S  be  directly  added  to  the  sodium  nitro- 
ferricyanide  as  it  causes  oxidation  of  the  sulphur.  To  test  for  ammonia 
the  gas  should  be  liberated  by  KOH  and  distilled  into  a  solution  of  H2S  ; 
and  this  solution  added  to  the  Na2Fe(NO)(CN)5 . 

c. — Ammonium  nitrite,  NH4NO2  ,  is  used  in  the  preparation  of  nitrogen 
(§235,  3);  ammonium  nitrate  in  the  preparation  of  nitrous  oxide,  N2O  , 
"  laughing-  gas"  (§237).  d. — Ammonium  phosphate,  as  a  reagent,  acts 
similarly  to  sodium  phosphate.  When  sodium  phosphate,  Na2HPO4  ,  is  used  to 
precipitate  metals  in  the  presence  of  ammonium  hydroxide,  a  double  phosphate 
of  the  metal  and  ammonium  is  frequently  formed  as  MnNH4PO4,  MgNH4PO4, 
etc.  By  some  chemists  microcosmic  salt,  NaNH4HPO4  ,  is  preferred  to  sodium 
phosphate,  Na,HPO4  ,  as  a  reagent. 

e. — When  ammonium  hydroxide  is  saturated  with  H2S ,  ammonium  sul- 
phide, (NH4)2S ,  is  formed.  Complete  saturation  is  indicated  by  the  failure 


§207,  6k.  AMMONIUM.  237 

to  precipitate  magnesium  salts,  that  is,  NH4OH  precipitates  magnesium 
salts  while  (NH4)2S  does  not.  Freshly  prepared  ammonium  sulphide  is 
colorless,  but  upon  standing  becomes  yellow  with  loss  of  ammonia  and 
formation  of  the  poly-sulphides,  (NH4)2SX .  The  yellow  poly-sulphide 
may  also  be  formed  by  dissolving  sulphur  in  the  normal  ammonium  sul- 
phide. As  a  precipitant  ammonium  sulphide  acts  similarly  to  the  fixed 
alkali  sulphides.  The  sulphides  of  Sb'"  and  Sn"  are  with  great  difficulty 
soluble  in  the  normal  ammonium  sulphide,  but  readily  soluble  in  the 
poly-sulphide.  Nickel  sulphide,  NiS,  is  insoluble  in  normal  ammonium 
sulphide  but  is  sparingly  soluble  in  the  yellow  poly-sulphide  (distinction 
from  cobalt).  (NH4)2S  gives  a  rich  purple  color  with  sodium  nitroferri- 
cyanide  (&).  Ammonium  sulphate  as  a  precipitating  reagent  acts  similar 
to  all  soluble  sulphates  (§205,  6e).  A  25  per  cent  solution  of  (NH4)2S04 
is  used  to  dissolve  CaS04  (§188,  5c)  (distinction  from  Ba  and  Sr). 

f. — Ammonium  chloride  is  much  used  as  a  reagent.  It  prevents  pre- 
cipitation of  the  salts  of  Mn  by  the  NH4OH ,  and  is  of  special  value  in  the 
precipitation  of  the  third  group  as  hydroxides  and  the  fourth  group  as 
sulphides  by  preventing  the  formation  of  soluble  colloidal  compounds. 
The  solubility  of  the  precipitates  of  the  carbonates  of  the  fifth  group  is 
slightly  increased  by  the  presence  of  ammonium  chloride;  i.  e.,  very  dilute 
solutions  of  barium  chloride  are  not  precipitated  by  ammonium  carbonate 
in  presence  of  a  large  excess  of  ammonium  chloride.  The  salts  of  mag- 
nesium are  not  precipitated  by  the  alkalis  or  by  the  alkali  carbonates  in 
presence  of  ammonium  chloride.  The  solubility  of  A1(OH)3  is  diminished 
by  the  presence  of  NH4C1  (§124,  6a,  and  §117). 

</,  //. — Similar  as  reagents  to  the  corresponding-  potassium  salts,  i. — Fliio- 
silicic  acid,  H2SiF6  ,  does  not  precipitate  ammonium  salts,  the  ammonium 
fluosilicate  being  very  soluble  in  water  (distinction  from  potassium),  j. — Chlor- 
platinic  acid,  H2PtCl6  ,  forms  with  ammonium  salts  the  yellow  ammonium 
chlorplatinate,  (NH^PtCle  ,  very  closely  resembling  the  potassium  salt  with 
the  same  reagent,  but  upon  ignition  only  the  spongy  metallic  platinum  is  left, 
i.  e.,  no  chloride  of  the  alkali  metal,  as  KC1 . 

k. — A  solution  of  potassium  mercuric  iodide,  K2HgI4,  containing  also 
potassium  hydroxide — Nessler's  test  * — produces  a  brown  precipitate  of 
nitrogen  dimercuric  iodide,  NHg2I,  dimercur-ammoiiium  iodide  (§58,  6a), 
soluble  by  excess  of  KI  and  by  HC1  ;  not  soluble  by  KBr  (distinction  from 
HgO): 

NH3  +  2HgI2  =  NHg2I  +  SHI 

NH<OH  +  2K2HgI4  -f  3KOH  =  NHg2I  +  7KI  +  4H2O 

*  This  reagent  may  be  prepared  as  follows  :  To  a  solution  of  mercuric  chloride  add  solution 
of  potassium  iodide  till  the  precipitate  is  nearly  all  redissolved  ;  then  add  solution  of  potassium 
hydroxide  sufficient  to  liberate  ammonia  from  ammonium  salts ;  leave  until  the  liquid  becomes 
gjear,  and  decent  from  any  remaining  sediment, 


AMMONIUM.  §207,  6Z, 

This  very  delicate  test  is -applicable  to  ammonium  hydroxide  or  salts; 
traces  forming  only  a  yellow  to  brown  coloration.  The  potassium  mercuric 
iodide,  "  Meyers^Keagent,"  alone,  precipitates  the  alkaloids  from  neutral 
or  acid  solutions,  but  does  not  precipitate  ammonium  salts  from  neutral 
or  acid  solutions.  Ammonium  hydroxide  in  alcoholic  solution  does  not 
give  a  precipitate  with  Nessler's  reagent,  but  from  this  solution  a  precipi- 
tate is  formed  with  HgCl2  (De  Koninck,  Z.,  1893,  32,  188). 

/. — Mercuric  chloride,  HgCl2 ,  forms,  in  solutions  of  ammonium  hy- 
droxide or  ammonium  carbonate,  the  "  white  precipitate "  of  nitrogen 
dihydrogen  mercuric  chloride,  NH2HgCl ,  or  mercur-ammonium  chloride. 
If  the  ammonium  is  in  a  salt,  not  carbonate,  it  is  changed  to  the  carbonate 
and  precipitated,  by  addition  of  mercuric  chloride  and  potassium  carbonate 
previously  mixed  in  solutions  (with  pure  water),  so  dilute  as  not  to  precipi- 
tate each  other  (yellow).  This  test  is  intensely  delicate,  revealing  the 
presence  of  ammonia  derived  from  the  air  by  water  and  many  substances 
(Wittstein,  Arch.  Pharm.,  1873,  203,  327). 

m. — Add  a  small  quantity  of  recently  precipitated  and  well-washed  silver 
chloride,  and,  if  it  does  not  dissolve  after  agitation,  then  add  a  little  potassium 
hydroxide  solution.  The  solution  of  the  Ag€l ,  before  the  addition  of  the  fixed 
alkali,  indicates  free  ammonia;  after  the  addition  of  the  fixed  alkali,  ammonium 
salt.  (Applicable  in  absence  of  thiosulphates,  iodides,  bromides  and  sulpho- 
cyanates.) 

n. — Sodium  phosphomolybdate  (§75,  Qd)  precipitates  ammonium  from  neutral 
or  acid  solutions;  also  precipitates  the  alkaloids,  even  from  very  dilute  solu- 
tions, and,  from  concentrated  solutions,  likewise  precipitates  K ,  B/b  and  Cs 
(all  the  fixed  alkalis  except  Na  and  Li). 

7.  Ignition. — Heat  vaporizes  the  carbonate,  and  the  haloid  salts  of  am- 
monium, undecomposed  (dissociated  but  reuniting1  upon  cooling) ;  decomposes 
the  nitrate  with  formation  of  nitrous  oxide  and  water,  and  the  phosphate  and 
berate  with  evolution  of  ammonia.  NH3  heated  to  780°  or  higher  is  dissociated 
into  N"  and  H  (Ramsay  and  Young,  J.  C.,  1884,  45,  88). 

8.  Detection. — As  ammonium  hydroxide  and  chloride  are  used  in  the 
regular  process  of  analysis,  the  original  solution  must  be  tested  for  the 
presence  or  absence  of  ammonium  compounds.     The  hydroxide  or  the 
carbonate  may  be  detected  by  the  odor  (1) ;  the  action  on  red  litmus  paper 
suspended  in  the  test-tube  above  the  heated  solution;  the  blue  color  im- 
parted to  paper  wet  with  copper  sulphate;  the  blackening  of  mercurous 
nitrate  paper;  and  if  in  considerable  quantity,  the  white  vapors  when 
brought  into  contact  with  the  vapors  of  volatile  acids.     In  combination 
as  salts  the  gas  is  liberated  by  the  fixed  alkali  hydroxides  or  carbonates 
(oxides  or  hydroxides  of  Ba ,  Sr ,  or  Ca  may  be  used)  and  distilled  into 
Nessler's  reagent,  or  collected  in  water  and  the  test  with  HgCl2  (6Z)  applied 
or  any  of  the  tests  for  ammonium  hydroxide. 

9.  Estimation. — Ammonium  salts  are  usually  estimated  by  distillation  into  a 
standard  acid,  from  a  solution  made  alkaline  with  KOH  ,  and  titration  of  the 
excess  of  the  acid  with  a  standard  NH4OH  solution,  using  tincture  of  cochineal 


§208, 5.  CAESIUM.  239 

as  an  indicator.     It  may  be  converted  into  the  chloride  and  precipitated  by 
PtCl4  and  weighed  as  the  double  platinum  salt. 

10.  Oxidation. — Ammonium  salts  in  solution,  treated  with  dtlorinc  gas,  gen- 
erate the  unstable  and  violently  explosive  "nitrogen  chloride"  (NCi3V)  (<i). 
The  same  product  is  liable  to  arise  from  solid  ammonium  salts  treated  with 
chlorine.  Gaseous  ammonia,  and  ammonium  hydroxide,  with  chlorine  gas, 
generate  free  nitrogen  (&),  a  little  ammonium  chlorate  being  formed  if  the 
ammonia  is  in  excess.  Hypochlorite-s  or  Jiypobromites  (or  chlorine  or  bromine 
dissolved  in  aqueous  alkali,  so  as  to  leave  an  alkaline  reaction)  liberate,  from 
dissolved  ammonium  salts,  all  of  their  nitrogen  (as  shown  in  the  second  equa- 
tion of  ?>);  the  measure  of  the  nitrogen  gas  being  a  means  of  quantitative 
estimation  of  ammonium.  With  iodine,  ammonium  iodide  and  the  explosive 
iodamides  (c)  are  produced;  or  under  certain  conditions  an  iodate  (d).  Ammo- 
nium hydroxide  is  liable  to  atmospheric  oxidation  to  ammonium  nitrite  and 
nitrate.  Permanganates  oxidize  to  nitrate  (e)  (Wanklyn  and  Gamgee,  J.  (7., 
1868,  21,  29).  In  presence  of  Cu  the  O  of  the  air  oxidizes  the  nitrogen  of 
ammonia  to  a  nitrite  (/)  (Berthelot  and  Saint-Gilles,  A.  Ch.,  1864,  (4),  1,  381). 
Ammonia  is  somewhat  readily  produced  from  nitric  acid  by  strong  reducing 
agents  (g).  It  is  formed  with  carbonic  anhydride,  in  a  water  solution  of 
cyanic  acid,  and,  more  slowly,  in  a  water  solution  of  hydrocyanic  acid.  It  is 
generated,  by  fixed  alkalis,  in  boiling  solution  of  cyanides  (ft) ;  also  in  boiling 
solutions  of  albuminoids  and  other  nitrogenous  organic  compounds,  this  forma- 
tion being  hastened  and  increased  by  addition  of  permanganate  (Wanklyn's 
process).  Fusion  with  fixed  alkalis  transforms  all  the  nitrogen  of  organic 
bodies  into  ammonia. 

(a)     NH4C1  +  3C12  =  NC13  +  4HC1 

(&)     8NH3  +  3C12  —  6NH4C1  +  N2 

2NH4C1  +  301,  =  8HC1  +  N2 

(c)  2NH3  +  I2  =  NH4I  +  NH2I 

(d)  6NH4OH  +  3I2  =  5NH4I  +  NH4I03  +  3H20 

(e)  6NH4OH  +  8HMnO4  =  3NH4NOS  +  8MnO(OH)a  +  5H2O 

(f)  12Cu  +  2NH3  +  902  =  12CuO  +  2HN02  +  2H2O 

(g)  3HN03  +  8A1  +  8KOH  =  8KA10,  +  3NH3  +  H2O 
(ft)     HCN  +  KOH  +  H20  =  NH3  +  KCHO2  (formate). 


§208.  Caesium.     Cs  =  132.81.    Valence  one. 

1.  Properties.— Specific  gravity,   1.88  at  15°    (Setterberg,  A.,   1882,  211,   100). 
Melting  point,   between   26°    and   27°.     It  is   quite   similar   to   the   other   alkali 
metals;    silver-white,    ductile,    very    soft    at    ordinary    temperature.     It    burns 
rapidly  when  heated  in  the  air,  and  takes  fire  when  thrown  on  water.     It  may 
be  kept  under  petroleum.     It  is  the  most  strongly  electro-positive  of  all  metals. 

2.  Occurrence.— Widely    distributed    but    in    small    quantities;    as    caesium 
aluminum  silicate  (mineral  castor  and  pollux)   (Pisani,  C.  r.,  1864,  58,  715);  in 
many  mineral  springs  (Miller,  C.  N.,  1864,  10,  181);  in  the  ash  of  certain  plants, 
tobacco,  tea,  etc. 

3.  Preparation. — By   electrolysis   of  a   mixture   of   CsCN  with   Ba(CN)2;   by 
ignition  of  CsOH  with  Al  in  a  nickel  retort  (Beketoff,  C.  C.,  1891,  (2),  450). 

4.  Oxide    and    Hydroxide. — An    oxide    has    not    yet    been    prepared.     The 
hydroxide,  CsOH  ,  is  a  gra^ash-white  solid,  very  deliquescent,  absorbs  CO2  from 
the  air;  dissolves  in  water  with  generation  of  much  heat,  forming  a  strongly 
caustic  solution. 

5.  Solubilities. — Caesium    dissolves    with    great    energy    in    water,    acids    or 
alcohol,   liberating   hydrogen  and  forming  the  hydroxide,   salts  or  alcoholate 
respectively.     The  hydroxide  is  soluble  in  water  and  alcohol.     The   salts   are 
all    quite    readily    soluble;     the    double   platinum   chloride,   Cs2PtCl4  ,    and    the 
acid  tartrate,  CsHC4H4O0  ,  being  least  soluble  and  used  in  preparation  of  the 
salts  free  from  the  other  alkali  metals. 


240  RUBIDIUM— LITHIUM.  §208,  6  . 

6.  Reactions. — In  all  its  reactions  similar  to  the  other  fixed  alkalis. 

7.  Ignition.     Caesium   sails  ceic;-    ine   non-luminous  flame   violet.     The   spec- 
trum gives  two  sharply  defined  lines,  Cs  a   and  Cs   ,1,  in  the  blue  and  a  third 
faint  line  in  the  orange-red   Us  t  ,  also  several  faint  lines  in  the  yellow  and 
green.     With  the  spectroscope  three  parts  of  CsCl  may  be  detected  in  presence 
of  300,000  to  400,000  parts  KC1  or  NaCl;  and  one  part  in  presence  of  1,500,000 
parts  LiCl  (Bunsen,  Pogg.,  1875,  155,  633). 

8.  Detection. — By  the  spectroscope  (7  and  §210,  7). 

9.  -Estimation. — (1)  As  the  double  platinum  chloride;  (2)  as  the  chloride  with 
RbCl ,  estimation  of  the  amount  of  Cl  and  calculation  of  the  relative  amounts 
of  the  metals;  (3)  as  the  sulphate  obtained  from  ignition  of  the  acid  tartrate 
and  treatment  with  H,S04  (Bunsen,  Poyg.,  18G3,  119,  1). 


§209.  Rubidium.     Rb  =  85.45.     Valence  one. 

1.  Properties. — Specific  gravitij,    1.52    (Bunsen,  A.,    1863,    125,   367).     Melting 
point,  38°  (Cir.  B.  of  S.,   1915);  at    -10°  soft  as  wax.     A  lustrous  silver-white 
metal  with  a  tinge  of  yellow,  oxidizes  rapidly  in  the  air,  developing  much  heat 
and  soon  igniting.     Volatile  as  a  blue  vapor  below  a  red  heat.     The  metal  does 
not  keep  well  under  petroleum,  but  is  best  preserved  in  an  atmosphere  of  hydrogen. 
Next  to  caesium  it  is  the  most  electro-positive  of  all  metals. 

2.  Occurrence. — Widely  distributed  in  small  quantities,  usually  with  caesium, 
and  frequently  with  the  other  alkali  metals,  always  in  combination.     None  of 
the  alkali  metals  can  occur  free  in  nature. 

3.  Preparation. — From  the  mother  liquor  obtained  in  the  preparation  of  Li 
salts   (Heintz,  J.  pr.,  1862,  87,  310):   (1)   By  ignition  of  the  acid  tartrate  with 
charcoal;    (2)    electrolysis    of    the    chloride;    (3)    by    ignition    with    Mg"   or    Al 
(Winkler,  B.,  1890,  23,  51;  Beketoff,  B.,  1888,  21,  c,  424). 

4.  Oxide  and  Hydroxide.- — The  oxide  Rb2O  has  not  been  with  certainty  pre- 
pared.    The  hydroxide,  RbOH  ,  is  formed  when  the  metal  is  decomposed   by 
water;  also  through  the  action  of  Ba(OH),  upon  Rb^S04  .     It  is  a  gray-white, 
brittle  mass,  melting  under  a  red  heat. 

5.  Solubilities. — The  metal  dissolves  in  cold  water,  in  acids  and  in  alcohol 
with   great   energy,  evolving  hydrogen.     The   hydroxide   is  readily   soluble   in 
water  with  generation  of  heat.     The  salts  are  all  quite  readily  soluble.     The 
acid  tartrate  is  about  eight  times  less  soluble  than  the  corresponding  Cs  salt. 
Among  the  less  soluble  salts  are  to  be  mentioned  the  perchlorate,  the  fluosili- 
cate,   the  double  platinum  chloride,  the   silicotungstate,   the  picrate,   and   the 
phosphomolj'bdate.     The  alum  is  less  soluble  than  the  corresponding  potassium 
alum. 

G.  Reactions. — Similar  to  the  other  fixed  alkalis. 

7.  Ignition. — The  salts  give  a  violet  color  to  the  flame.     The  spectrum  gives 
two  characteristic  lines  in  the  violet,  Rb  a- and  Rb  ,,?;  two  less  intensive  in  the 
outer  red,  Rb  >  and  Rb  J;  a  fifth  Rbtin  the  orange;  and  many  faint  lines  in  the 
orange,  yellow  and  green.     As  small  a  quantity  as  0.0000002  gram  of  RbCl  can 
be  detected  (Bunsen,  I.e.). 

8.  Detection. — By  the  spectroscope  (7  and  §210,  7). 

9.  Estimation.—  (1)  By  weighing  with  CsCl  as  the  chlorides,  determining  the 
amount  of  Cl  and  calculating  the  proportion  of  the  metals;   (2)  as  the  double 
platinum  chloride. 

§210.  Lithium.    Li  =  6.94.     Valence  one. 

1  Frcpsitie;. — Specific  gravity,  0.5936,  the  lightest  of  all  known  solid  bodies 
(Bunsen  and  Matthiessen,  A.,  1855,  94,  107).  Melting  point,  186°  (Cir.  B.  of 
S.  1915);  does  not  vaporize  at  a  red  heat.  It  is  a  silver-white  metal  with  a 
grayish  tinge;  harder  than  K  or  Na,  but  softer  than  Pb  ,  Ca  or  Sr;  it  is  tough 
and  may  be  drawn  into  wire  and  rolled  into  sheets.  It  is  more  electro-positive 
than  the  alkaline  earth  metals,  but  less  electro-positive  than  K  or  Na  .  The 
pure  metal  is  quite  similar  in  appearance  and  in  its  chemical  properties  to  K 


§210, 


LITHIUM.  241 


and  Na ,  but  does  not  react  so  violently  as  those  metals.  It  does  not  ignite 
in  the  air  until  heated  to  '200°,  and  then  burns  quietly  with  a  very  intense  white 
light.  It  also  burns  with  vivid  incandescence  in  Cl  ,  Br  ,  I  ,  O  ,  S  and  dry  CO,.  . 
It  decomposes  water  readily,  forming  LiOH  and  H,  but -not  with  combustion 
of  the  hydrogen  or  ignition  of  the  metal. 

2.  Occurrence. — It  is  a  sparingly  but  widely  distributed   metal,  usually   pre- 
pared   from    lepidolite,    triphylite,    or    petalite.     Traces    are    found    in    a    great 
many  minerals,  in  mineral  spring's,  and  in  the  leaves  and  ashes  of  many  plants; 
e.  g.,  coffee,  tobacco  and  sugar-cane. 

3.  Preparation. — It    is   prepared   pure    only   by   electrolysis,    usually    of    the 
chloride.     A  larger  yield  is  obtained  by  mixing"  the  LiCl  with  NH4C1  or  KC1 
(Giintz,  C.  r.,  1S9.'5,   117,  7!!2).     The  metal  is  also  obtained  by  ignition  of  the 
carbonate  with  Mg1 ,  but  the  metal  is  at  once  vaporized  and  oxidized. 

4.  Oxide  and  Hydroxide. — It  forms  one  oxide,  Li^O  ,  by  heating  the  metal 
in  oxygen  or  dry  air;  cheaper  by  the  action  of  heat  upon  the  nitrate.     The 
corresponding  hydroxide,   LiOH  ,   is   made   by   the   action   of  water  upon   the 
metal  or  its  oxide:  cheaper  by  heating  the  carbonate  with  calcium  hydroxide. 

5.  Solubilities. — The    metal    is    readily    soluble    in    water   with    evolution    of 
hydrogen,   forming  the   hydroxide;   soluble  in  acids  with   formation   of   salts. 
The  oxide,  LLO  ,  dissolves  in  water,  forming  the  hydroxide.     The  most  of  the 
lithium    salts    are    soluble    in   water.     A    number    of    the   salts,    including   the 
chloride    and   chlorate,    are   very   deliquescent.     The   hydroxide,    carbonate    and 
phosphate  are  less  soluble  in  water  than  the  corresponding  compounds  of  the 
other  alkali  metals.     In  this  respect  lithium  shows  an  approach  to  the  alkaline 
earth  metals.     LiOH  is  soluble  in  14.5  parts  water  at  20°  (Dittmar,  J.  Soc.  Ind., 
1888,  7,  730);  Li,C03  in  75  parts  at  20°;  Li3PO4  in  2539  parts  pure  water  and 
3920  parts  ammoniacal  water,  more  soluble  in  a  solution  of  NH4C1  than  in 
pure  water   (Mayer,   A.,    i66o,   i,8?    193).     Lithium  chloroplatinate  and  lithium 
picrate  are  very  soluble  in  water  (Richard  Z.,  40,  383). 

G.  Reactions. — Lithium  salts  in  general  react  similar  to  the  corresponding 
potassium  and  sodium  salts.  They  are  as  a  rule  more  fusible  and  more  easily 
decomposed  upon  fusion.  Soluble  phosphates  precipitate  lithium  phosphate, 
more  soluble  in  NH,C1  solution  than  in  pure  water  (distinction  from  mag- 
nesium). In  dilute  solutions  the  phosphate  is  not  precipitated  until  the  solu- 
tion is  boiled.  The  delicacy  of  the  test  is  increased  by  the  addition  of  NaOH, 
forming  a  double  phosphate  of  N.a  and  Li  (Rammelsberg,  A.  Ch.,  1818,  (2),  7, 
157).  The  phosphate  dissolved  in  HC1  is  not  at  once  precipitated  by  neutraliz- 
ing with  NH4OH  (distinction  from  the  alkaline  earth  metals).  Nitrophenic 
acid  forms  a  yellow  precipitate,  not  easily  soluble  in  water. 

7.  Ignition. — Compounds  of  lithium  impart  to  the  flame  a  carmine-re^  color, 
obscured   by   sodium,   but    not   by   small    quantities   of   potassium   compounds. 
Blue  glass,  just  thick  enough  to  cut  off  the  yellow  light  of  sodium,  transmits 
the  fed  light  of  lithium;  but  the  latter  is  intercepted  by  a  thicker  part  of  the 
blue  prism,  or  by  several  plates  of  blue  glass.     The  spectrum  of  lithium  con- 
sists of  a  bright  red   band,  Li  a,   and  a  faint  orange  line,  Li  /?.     The  color 
tests  have  an  intensity  intermediate  between  those  of  sodium  and  potassium. 

8.  Detection. — #.//  the  sprctroscope. — To  the  dry  chlorides  of  the  alkali  metals 
a  few  drops  of  HC1  are  added  and  the  mass  extracted  with  90  per  cent  alcohol. 
The  solution  contains  all  the  rare  alkalis  and  some  Na  and  K  .     Evaporate  to 
dryness,  dissolve  in  a  small  amount  of  water  and  precipitate  with  platinum 
chloride.     The  double  platinum  and  potassium  chloride  is  more  soluble  than 
the  corresponding  salt  of  Rb  and  Cs  .     Boil  repeatedly  with  small  portions  of 
water  to   remove  the  potassium,   and   frequently   examine  the  residue  by   the 
spectroscope  as  follows:  Wrap  a  small  amount  of  the  precipitate  in  a  moistened 
filter  paper,  then  in  a  platinum  wire  and   carefully  char.     After  charring  is 
complete,  ignite  before  the  spectroscope.     The  K  spectrum  grows  fainter,  that 
of  Rb  and  Cs  appear. 

Evaporate  to  dryness  the  filtrate  from  the  precipitate  of  the  platinum  double 
salts,  add  oxalic  acid  and  ignite,  moisten  with  HC1,  evaporate  and  extract  with 
absolute  alcohol  and  ether.  Upon  evaporation  of  the  extract  LiCl  is  obtained, 
almost  pure.  Test  with  the  spectroscope  and  by  forming  the  insoluble  phos- 
phate. 


242  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES.  §210,9, 

9.  Estimation. — After  separation  from  other  elements  it  may  be  weighed  as 
a  sulphate,  carbonate  or  phosphate,  Li3PO4  .  It  may  also  be  estimated  by  the 
comparative  intensity  of  the  lines  in  the  spectroscope  (Bell,  Am.,  1886,  7,  35). 


DIRECTIONS  FOR  THE  ANALYSIS  OF  THE  METALS  OF  THE  ALKALI  GROUP 

(SIXTH  GROUP). 

§211.  If  the  material  is  found  not  to  contain  magnesium,  the  clear 
filtrate  from  the  carbonates  of  Ba ,  Sr ,  and  Ca ,  after  testing  for  traces 
with  (NH4)2S04  and  (NH4)2C204  (§193),  may  at  once  be  tested  for  the  pres- 
ence of  potassium  and  sodium.  If  magnesium  be  present  it  should  be 
removed  in  order  to  test  for  small  amounts  of  sodium.  Potassium  and 
large  amounts  of  sodium  may  be  readily  detected  in  the  presence  of  mag- 
nesium. It  is  evident  that  the  magnesium  must  not  be  removed  by  the 
usual  reagent  used  to  detect  the  presence  of  that  element,  i.  e.  Na2HP04 . 
It  is  recommended  by  many  to  use  ammonium  phosphate.  (NH4)2HP04 . 
This  reagent  removes  the  magnesium,  and  permits  the  application  of  the 
flame  test  for  the  fixed  alkalis;  but  the  presence  of  the  phosphate  obstructs 
the  gravimetric  determination  of  the  alkalis.  The  phosphate  may  be 
removed  by  lead  acetate  and  the  excess  of  the  lead  by  hydrogen  sulphide. 

§212.  As  a  better  method  it  is  directed  to  evaporate  the  filtrate  con- 
taining the  magnesium  and  the  alkalis  to  dryness,  ignite  gently  to  remove 
the  ammonium  salts.  Dissolve  the  residue  in  water  and  add  Ba(OH)2  to 
precipitate  the  magnesium  as  Mg(OH)2  (§§177  and  182).  After  filtration, 
the  excess  of  barium  in  the  filtrate  is  removed  by  H2S04 ,  and  the  filtrate 
from  the  barium  sulphate  is  ready  to  be  tested  for  the  fixed  alkalis  by  the 
flame  test  or  by  gravimetric  methods  as  may  be  desired.  The  presence  of 
sodium  obscures  the  flame  reaction  for  potassium,  but  the  introduction 
of  a  cobalt  glass  (§132,  7)  or  an  indigo  prism  cuts  out  the  sodium  flame 
and  allows  the  violet  potassium  flame  to  be  seen.  Study  6,  7,  8,  and  9  of 
§§205  and  206. 

§213.  The  free  use  of  ammonium  salts  during  the  process  of  analysis 
makes  it  necessary  that  the  testing  for  ammonium  be  done  in  the  original 
solution  or  in  the  filtrate  from  the  Tin  and  Copper  Group. 

Add  an  excess  of  KOH  or  NaOH  to  the  solution  and  warm  gently.  Notice 
the  odor  (§207,  1).  Suspend  a  piece  of  moistened  red  litmus  paper  in 
the  test-tube;  in  the  presence  of  ammonia  it  will  be  changed  from  red 
to  blue  color.  To  detect  the  presence  of  small  amounts  of  ammonium 
salts,  heat  the  strongly  alkaline  mixture  nearly  to  boiling  and  pass  the 
evolved  gas  into  water.  Test  this  solution  (ammonium  hydroxide)  with 
Nessler's  Reagent  (§207,  6fc)  or  by  the  precipitation  with  HgCl2  (§207,  6/). 
Study  §207,  6,  7,  8,  and  9. 

§214.  The  rare  metals  of  the  Alkali  Group:  lithium,  rubidium,  and 


§215.  DIRECTIONS  FOR  ANALYSIS  WITH  NOTES.  243 

caesium,  are  rarely  met  with  in  the  ordinary  analyses.  If  their  presence 
is  suspected  they  are  tested  for  and  detected  by  the  spectroscope  (7,  §§208, 
209  and  210). 

§215.  Lithium,  because  of  the  insolubility  of  its  phosphate  (§210,  5c), 
interferes  with  the  detection  of  magnesium.  If  the  nitrate  after  the 
removal  of  barium,  strontium,  and  calcium  be  evaporated  to  dryness  and 
gently  ignited  to  remove  all  ammonium  salts,  the  residue,  dissolved  in 
water  and  treated  with  an  excess  of  barium  hydroxide,  will  give  a  precipi- 
tate of  the  magnesium  as  the  hydroxide,  leaving  the  lithium  in  solution. 
The  barium  hydroxide  precipitate  may  be  tested  for  magnesium  and  from 
the  nitrate  the  excess  of  barium  hydroxide  may  be  removed  by  sulphuric 
acid  before  testing  for  the  alkali  metals. 


PART  III -THE  NON-METALS. 


§216.  BALANCING  EQUATIONS  IN  OXIDATION  AND  EEDUCTION. 

Oxidation  and  reduction  always  involves  a  change  in  valence.  When  the 
valence  of  an  electropositive  element  is  increased  it  is  said  to  be  oxidized, 
and  conversely,  when  its  valence  is  reduced,  reduction  has  taken  place.  It 
is  believed  that  each  bond  or  valence  is  produced  by  the  presence  of  a  unit 
charge  of  electricity  on  the  atom  or  ion.  A  ferrous  ion  may  be  represented 
as  Fe++  while  a  ferric  ion  would  be  Fe+  +  +  and  the  oxidation  or  reduction 
of  iron  would  therefore  consist  in  adding  to  or  subtracting  a  unit 
charge  of  electricity  from  the  atom.  Similarly  the  valence  of  negative  ele- 
ments is  proportional  to  the  number  of  negative  charges  of  electricity  on 
their  atoms.  It  is  assumed  that  during  oxidation  and  reduction  unit 
charges  of  negative  electricity  carried  by  small  corpuscles  or  electrons  pass 
from  one  atom  to  the  other.  The  positively  charged  masses  of  atoms  are 
very  much  larger  than  the  negative  electrons.  When  an  atom  loses  a  nega- 
tive electron,  its  positive  charge  is  relatively  increased  and  it  is  oxidized; 
when  it  gains  a  negative  electron  it  is  reduced. 

The  valence  of  an  electro-negative  element  would  therefore  be  increased 
when  it  is  reduced  and  reduced  when  it  is  oxidized.  When  an  element 
can  pass  from  the  positive  to  the  negative  condition  there  may  be  no  change 
in  valence  during  reduction  or  oxidation.  2NfH^  -f  30=  =N^"  +  +03=  + 
3H20. 

The  metals  in  salts  generally  act  as  electro-positive  elements  while  the 
acid  elements  or  radicles  are  electro-negative.  The  latter  therefore  may 
act  as  oxidizing  agents  towards  the  former.  In  general  in  a  reaction  in- 
volving oxidation  and  reduction  one  element  is  oxidized  and  another  is 
reduced  and  the  gain  in  valence  of  one  element  is  exactly  equal  to  the  loss 
in  valence  of  the  other  element.  This  is  a  necessary  consequence  of  the 
transfer  of  negative  electrons  from  one  element  to  the  other. 

Statement  of  Bonds  in  Plus  and  Minus  Numbers,*  according  to  chemical 
polarity,  positive  and  negative  (see  §3  footnote). 

A  bond,  that  is  a  unite  of  active  valence,  is  either  a  plus  one  or  a  minus 
one.  The  formula  of  a  molecule  of  hydrochloric  acid  is  stated,  H+IC1~I, 

*  O.  C.  Johnson,  C.  N.,  1880,  42,  51.  See  also  Ostwald,  Grundr.  allg.Chem.,  3te  Aufl.,  1899, 
S.  439. 


§216.      BALANCING   EQUATIONS    IN   OXIDATION   AND    REDUCTION.       245 

that  of  water,  (H+I)20~n.  (The  plus  sign  is  understood  when  no  sign  is 
written  before  the  valence  number.) 

Plus  and  minus  bonds  are  represented  as  positive  and  negative  quan- 
tities. In  the  formula  of  hydrochloric  acid,  as  above,  the  difference 
between  the  polarity  of  the  hydrogen  atom  and  that  of  the  chlorine  atom 
is  stated  as  a  difference  of  two. 

In  any  compound  the  sum  of  the  plus  bonds  and  the  minus  bonds  of  the 
atoms  forming  a  molecule  is  zero. 

Free  elements,  not  having  active  valence,  have  zero  bonds  in  this 
notation.* 

The  Oxidation  of  any  element  is  shown  by  an  increase,  and  its  Eeduction 
by  a  decrease,  in  the  sum  of  its  bonds. 

When  one  substance  reduces  another  the  element  which  is  reduced 
loses  as  many  bonds  as  are  gained  by  the  element  which  is  oxidized. 

It  is  evident  that,  changes  in  valence  being  reciprocal  in  oxidation  and 
reduction,  there  is  no  gain  or  loss  in  the  sum  of  the  bonds  of  two  elements 
which  act  upon  each  other. 

The  use  of  this  notation  is  illustrated  in  the  following  equations  : 
3SnCl,  +  H,S03  +  6HC1  =  3SnCl4  +  H2S  +  3H20 

In  this  equation  the  three  atoms  of  tin  gain  six  bonds;  the  bonds  of  the 
sulphur  in  the  H2SO,  have  then  been  diminished  by  six;  that  is,  it  has 
given  up  six  bonds  to  the  tin,  and  having  only  four  in  the  first  place  must 
now  have  minus  two  (4  -6  =  -2). 

The  valence  of  the  acid  element  in  an  acid  may  always  be  found  from 
the  anhydride.  In  this  case  we  have:  H2S03=S02+H20  ,  the  valence  of 
the  sulphur  in  S02  being  4. 


3Sn  C'2  +  FI03  +  6HC1  =  3Sn    C14  +HI  +  3H2O 

Here  also  the  three  atoms  of  tin  gain  six  bonds,  and  these  are  furnished 
by  the  iodine  of  the  HIO.,  .  It  has  five  in  the  first  place,  and  being 
diminished  by  six,  has  one  negative  bond  remaining  (5  -6=-]).  [In 
other  words,  unless  we  deny  that  iodine  has  five  bonds  in  HI03  ,  we  must 
admit  that  it  has  one  negative  bond  in  HI  (written  HT~').] 

8HMn04  +  5AsH3  +  8H2S04  =  5H3AsO4  +  8MnSO4  +  12H20 
In  this  equation  eight  atoms  of  manganese  in  the  first  member  have  56 
bonds,  and  a  like  amount  in  the  second  member  has  only  16,  losing  40, 
and  this  40  has  been  gained  by  the  five  atoms  of  arsenic.     They  now  have 

*  If  there  is  polarity  in  the  union  of  like  atoms  with  each  other  in  forming  an  elemental 
molecule,  the  sum  must  be  zero,  as  in  the  formation  of  the  molecules  of  compounds. 


246  BALANCING   OF   EQUATIONS,  §217. 

25,  after  gaining  40.  They  must  then  have  had  — 15  in  the  first  place 
(25  —  40  — -15).  That  is,  the  atom  of  arsenic  in  arsenous  hydride  has 
-3  bonds  (As-'"H3). 

SnCl2  +  HgCL  —  Hg  +  SnCl4 

This  equation  illustrates  the  statement  that  free  elements  have  no 
bonds.  The  tin  gains  two  bonds,  and  these  two  bonds  are  taken  from  the 
mercury  in  the  HgCl2 . 


§217.  Rule  for  Balancing  Equations. 

The  number  of  oxidation  bonds  which  any  element  has  is  determined 
by  the  following  rules : 

a.  Hydrogen  has  always  one  positive  bond. 

b.  Oxygen  has  always  two  negative  bonds. 

c.  Free  elements  have  no  bonds. 

d.  The  sum  of  the  bonds  of  any  compound  is  zero. 

e.  In  salts  the  bond  of  the  metal  is  always  positive. 

/.  In  acids  and  in  salts  the  acid  radical  has  always  negative  bonds. 

Thus,  the  bond  of  free  Pb  is  zero,  but  in  PbCl2  the  lead  has  two  posi- 
tive bonds,  and  each  atom  of  chlorine  has  one  negative  bond. 

In  Bi2S3 ,  each  atom  of  Bi  has  three  positive  bonds  (e),  and  each  atom  of 
S  has  two  negative  bonds  (/). 

In  the  following  salts,  etc.,  the  bond  of  each  element  is  marked  above, 
with  its  proper  sign,  plus  being  understood  if  no  sign  is  given.  Then  fol- 
lows the  equation  in  full,  the  bonds  of  each  atom  being  multiplied  by  the 
number  of  atoms,  and  all  being  added,  the  sum  is  seen  to  be  zero. 

Hg"(NvO-"a)2.2  +  10  —  12  =  0 

Bi'"2(SviO-"4)3.6  +  18  —  24  =  0 

Ba"(MnViiO-"4)2.2  +  14  —  16  =  0 

Fe'"(NvO-"3)3.3  +  15  —  18  =  0 

As'"2S-"3.6  —  6  =  0 

If  the  above  is  understood,  the  rule  for  balancing  equations  is  easily 
explained. 

The  number  of  bonds  changed  in  one  molecule  of  each  shows  the  number 


§218,  4-  BALANCING   OF   EQUATIONS.  247 

of  the  molecules  of  the  other  which  must  be  taken,  the  words  each  and 
other  referring  to  the  oxidizing  and  reducing  agents. 


§218.  A  few  equations  will  illustrate  the  application  of  the  rule. 

(1)  3As4  +  20HN03  +  8H,O  =  12H3As04  +  20NO 

The  arsenic  in  one  molecule  gains  20  bonds,  therefore  20  molecules  of  HN03 
are  taken.  The  nitrogen  loses  three  bonds,  therefore  three  molecules  of  As4 
are  taken.  The  valence  of  the  nitrogen  and  arsenic  may  be  found  from  their 
anhydrides: 

2HNO3  =  H2O  +  N2O5     and     1  H3AsO4  =  3H2O  +  As2O5 

Equations  of  this  kind  may  also  be  balanced  by  considering  that  the  arsenic 
must  be  oxidized  to  the  pentoxide,  one  molecule  of  arsenic  requiring  ten  atoms 
of  oxygen.  Two  molecules  of  nitric  acid  furnish  three  atoms  of  oxygen,  as 
follows : 

i'HNOs  =  H2O  +  2NO  +  3O 

Three  molecules  of  arsenic  must,  therefore,  be  taken,  requiring  thirty  atoms 
of  oxygen,  which  will  be  furnished  by  20  molecules  of  nitric  acid.  To  convert 
a  molecule  of  arsenic  pentoxide  into  arsenic  acid  requires  3  molecules  of  water. 
The  3  molecules  of  arsenic  will  therefore  require  18  molecules  of  water.  As 
the  20  molecules  of  nitric  acid  furnish  10  molecules  of  water,  8  more  must  be 
added. 

(2)  6Sb  +  10HN03  =  3Sb205  +  10NO  +  5H2O 

The  antimony  gains  five  bonds,  therefore  five  molecules  of  HN03  would  be 
taken,  and  since  the  nitrogen  loses  three  bonds,  three  of  antimony  would  be 
taken,  but  since  we  cannot  write  Sb2O5  with  an  odd  number  of  atoms  of 
antimony,  we  double  the  ratio  and  take  six  and  ten. 

(3)  3H2S  +  8HN03  =  3H2S04  +  8NO  +  4H20 

The  S  in  the  first  member  has  2  negative  bonds  (a  and  d);  in  the  second 
member  it  has  6  positive,  gaining  8  bonds;  hence  8  molecules  of  HNO3  must 
be  taken.  The  nitrogen  in  the  first  member  has  five  bonds,  and  in  the  second 
it  has  two.  The  difference  is  three,  therefore  just  three  molecules  of  H2S 
must  be  taken. 

Further,  the  reaction  may  be  explained  as  follows: 

The  sulphur  in  the  first  member  has  two  bonds  (valence  of  two),  but  nega- 
tive because  combined  with  hydrogen  (two  atoms)  to  form  a.  definite  com- 
pound; in  the  second  member  it  has  six  bonds  (valence  of  six),  but  positive 

because  combined  with  oxygen  (SO3  or  JJQ  H  S  —  O^'  The  valence  of  the 
hydrogen  does  not  change  and  hence  in  the  reaction  one  molecule  of  H2S 
gains  eight  bonds.  The  nitrogen  in  the  first  member  has  five  bonds  (valence 

of  five),  but  positive  because  combined  with  oxygen  (N2O5  or  H  —  0  —  N~ ? ) ; 

in  the  second  member  it  has  two  bonds,  still  positive  because  combined  with 
oxygen.  The  valence  of  the  hydrogen  and  oxygen  does  not  change,  hence  in 
the  reaction  one  molecule  of  HNO.,  loses  three  bonds.  Now  the  number  of 
bonds  gained  by  the  HL,S  (8)  must  equal  the  bonds  lost  by  the  HN03  (3). 
The  least  common  multiple,  twenty-four,  indicates  the  least  possible  total 
change  of  valence  for  en  oh  compound;  this  requires  that  three  molecules  of 
H,S  and  eight  of  HNO3  be  taken,  giving  for  the  products  three  molecules  of 
ELSO*  and  eight  of  NO  with  four  of  water  to  complete  the  equation. 

(4)  3Sb2S8  +  28HN03  =  3Sb205  +  9H2S04  +  28NO  -f  5H2O 

In  this  case,  both  the  Sb  and  the  S  in  the  molecule  gain  bonds,  and  must  be 


248  BALANCING  OF  EQUATIONS.  §218,  o. 

considered.  It  is  plain  (from  d  and  e)  that  each  atom  of  Sb  gains  2  bonds,  and 
the  two  in  the  molecule  will  gain  4. 

The  S  in  Sb2S3  has  2  negathe  bonds,  and  in  the  second  member  (in  H2SO4) 
it  has  6  positive  bonds,  a  gain  of  8.  The  three  atoms  in  the  molecule  will  gain 
three  times  eight,  or  24  bonds;  to  this  add  the  4  which  the  Sb  has  gained,  and 
we  have  28  bonds  gained  by  one  molecule  of  Sb2S3;  hence  28  molecules  of  HNO3 
must  be  taken.  We  take  3  of  Sb2S3  for  reasons  explained  in  the  first  equation. 

Further  explain  as  follows:  In  this  case  both  the  Sb  and  the  S  gain  in 
valence  (oxidized).  Each  atom  of  antimony  gains  two  bonds,  a  total  gain  of 
four.  Each  atom  of  sulphur  gains  eight,  a  total  gain  of  twenty-four;  or  a 
gain  for  one  molecule  of  Sb2S3  of  twenty-eight  bonds.  As  in  the  previous 
illustration,  the  nitrogen  loses  three  bonds.  The  least  common  multiple, 
eighty-four,  indicates  that  for  the  reaction  each  compound  must  undergo  a 
change  of  at  least  eighty-four  bonds.  This  requires  for  the  Sb2S3  three  mole- 
cules, and  for  the  HN03  twenty-eight  molecules.  The  products  are  as  indicated 
in  the  equation. 

(5)  2Ag3As04  +  HZn  +  11H2S04  =  2AsH3  +  GAg1  +  llZnS04  +  8H2O 

The  silver  loses  three  bonds,  and  the  arsenic  in  changing  from  plus  five  to 
minus  three  loses  eight  bonds;  this  added  to  the  three  that  the  silver  loses 
makes  eleven,  therefore  eleven  atoms  of  zinc  are  taken,  and  since  the  zinc  gains 
two,  two  molecules  of  silver  arsenate  are  taken. 

(6)  2MnO  +  :Pb304  +  COHN08  =  2HMn01  +  15Pb(N03),  -}-  14H20 

The  manganese  gains  five  bonds,  therefore  five  molecules  of  Pb304  are  taken. 
The  three  atoms  of  lead  in  one  molecule  of  Pb3O4  have  in  all  eight  bonds,  but 
a  like  amount  has  only  six  in  the  second  member,  being  a  loss  of  two,  there- 
fore two  molecules  of  MnO  are  taken. 

(7)  2MnBr2  +  7Pb02  +  14HN03  =  2HMn04  +  2Br2  +  7Pb(N03)2  +  GH20 

The  manganese  gains  five  bonds  and  the  bromine  gains  one,  the  two  atoms 
gaining  two,  adding  this  to  the  five  that  the  manganese  gains  makes  a  total 
gain  of  seven  bonds,  therefore  seven  of  Pb02  are  taken.  The  lead  loses  two, 
therefore  two  of  MnBr,  are  taken. 

(8)  MnS  +  4KNO3  +  K2C03  ,  fusion  =  K2MnO4  +  K2S04  +  4NO  +  K2C03 

The  manganese  gains  four  bonds  and  the  sulphur  eight,  making  twelve; 
therefore  twelve  of  KNO3  would  be  taken,  and  since  the  nitrogen  loses  three 
bonds,  three  of  MnS  would  be  taken,  but  since  three  is  to  twelve  as  one  is  to 
four,  the  latter  amounts  are  taken. 

(9)  2Cr(OH)3  +  3Mn(N03)2  +  5K2CO3  ,  fusion  =  2K2Cr04  + 

3K,MnO4  +  6NO  +  5C02  +  3H2O 

This  is  a  peculiar  and  instructive  equation.  The  nitrogen  loses  six  bonds,  but 
since  the  manganese  in  the  same  molecule  gains  four,  the  total  loss  is  only  two, 
therefore  two  of  Cr(OH)3  are  taken.  The  chromium  gains  three,  therefore 
three  of  Mn(N03)2  are  taken. 

(10)  3Ag  +  4HN03  =  3Ag-N03  +  NO  +  2H20 

The  rule  here  calls  for  three  of  silver  and  one  of  nitric  acid,  but  three  more 
of  unreduced  nitric  acid  are  needed  to  combine  with  the  silver,  making  four 
in  all. 

(11)  2FeI2  +  6H2S04  ,  cone.,  hot  =  Fe2(S04)3  +  3S02  +  2I2  +  6H20 

The  rule  here  calls  for  two  of  FeI2  and  three  of  H2S04  ,  but  three  more  of 
H2SO4  that  are  not  reduced  are  needed  to  combine  with  the  iron,  making  six 
in  all. 

(12)  3HN08  +  8A1  +  8KOH  =  3NH3  +  8KA102  +  H20 

The  nitrogen  has  five  bonds  in  HN03  ,  and  in  NH3  it  has  minus  three, 
losing  eight,  therefore  eight  of  aluminum  are  taken.  The  aluminum  gains 
three,  therefore  three  of  HNO3  are  taken. 


§218,  15.  BALANCING  OF  EQUATIONS.  249 


(13)  3BiON03  +  11A1  +  11KOH  =  3Bi  +  3NH3  +  11KA1O2  +  H20 

The  bismuth  loses  three  bonds  and  the  nitrogen  loses  eight,  therefore  eleven 
of  aluminum  are  taken;  the  aluminum  gains  three,  therefore  three  of  the 
BiONO3  are  taken.  ' 

(14)  Mn(X  +  4HC1  =  MnCL  +  CL  +  2H20 

The  manganese  loses  two  bonds  and  the  chlorine  gains  one,  but  two  more  of 
unoxidized  HC1  are  needed  to  combine  with  the  manganese,  hence  four  are 
taken. 

(/5)     2CrI3  +  G4KOH  +  27CL  =  2K2Cr04  +  GKI04  +  54KC1  +  32H2O 

The  chromium  gains  three  bonds  and  the  iodine  (in  the  molecule)  gains 
twenty-four,  therefore  twenty-seven  of  C12  are  taken  and  the  C12  loses  two, 
therefore  two  of  CrI3  are  taken. 

This  rule  holds  good  in  organic  chemistry  when  all  the  products  of  the 
reactions  are  known,  as  the  following  examples  will  illustrate: 

CH4  C-4H'4.  —  4+4  =  0 

CH3C1  C-3  +  'H',Cl-/.  —3+1  +  3-1  =  0 

CH2Cla  C-2  +  2H'2Cl-'a.  —2+2+2-2  =  0 

CHOI,  C-1  +  8H'Cl-'3.  -1+3+1-3  =  0 

CCL  C4Cl-'4.  4—4  =  0 

HC2H8Oa  H'(Ca)  +  3-3H'3O-aa  .         1  +  3-3  +  3  —  4  =  0 

CaH6O  (C2)1-5H/6O-2.  1  —  5  +  6  —  2  =  0 

CSH8O8  (Cs)  -5  +  3  H'8O  -23  •  —5  +  3  +  8  —  6  =  0 

C6H1208  (C6)-7  +  7H'120-V  -7+7+12-12  =  0 

(1)  CH4  +  4C12  =  CC14  +  4HC1 

The  carbon  is  oxidized  by  the  chlorine  from  negative  four  to  positive  four, 
a  polarity  change  of  eight  units,  hence  take  eight  molecules  of  chlorine;  each 
molecule  of  chlorine  loses  two  bonds,  take  two  molecules  of  methane.  Two  is 
to  eight  as  one  is  to  four. 

(2)  3C2H60  +  2K2Cr20T  +  8H2S04  =  3HC2H3O2  +  2K2S04  + 

2Cr2(S04),  +  11H20 

The  carbon  of  the  alcohol  while  possessing  a  valence  of  eight,  has  an  oxida- 
tion valence  of  but  four  (minus  four  bonds)  ;  in  the  acetic  acid  the  two  atoms 
of  carbon  have  zero  bonds,  that  is,  the  combinations  with  negative  affinity 
exactly  equal  the  combinations  with  positive  affinity;  therefore  take  four 
molecules  of  the  potassium  dichromate.  The  two  atoms  of  the  chromium  lose 
six  bonds,  take  six  molecules  of  the  alcohol.  Six  is  to  four  as  three  to  two. 
Eight  molecules  of  sulphuric  acid  are  necessary  to  combine  with  the  potassium 
and  the  chromium. 

(3)  3C,H803  +  14HN03  =  9C02  +  14NO  +  19H.O 

The  three  atoms  of  the  carbon  in  the  glycerine  have  minus  two  bonds  (the 
negative  affinity  is  two  more  than  the  positive  affinity),  and  in  the  C02  a  like 
amount  has  twelve  bonds,  a  gain  of  fourteen.  The  nitrogen  loses  three  bonds. 

(4)  C0H1206  +  12H2S04  =  GC02  +  12S02  +  18H.O 

The  carbon  in  the  dextrose  has  zero  bonds  (equal  positive  and  negative 
affinity  combinations)  and  gains  twenty-four  bonds,  while  the  sulphur  loses 
two  bonds.  The  lower  ratio  is  one  to  twelve. 

For  convenience  of  reference  the  non-metallic  elements  will  be  de- 
scribed in  the  order  of  their  atomic  weights;  and  the  acids  in  the  order 
of  the  degree  of  oxidation  of  the  characteristic  element,  e.  g.,  N  before  S  , 
HC1  before  HC10  ,  HC103  before  HC104  ,  etc.. 


HYDROGEN.  §219 

§219.  Hydrogen.     H=  1.008  .     Valence  one. 

1.  Properties. — An  odorless,   tasteless  gas.     It  is   the  lightest  body   known: 
One  litre  at  0°,  760  mm.  atmospheric  pressure,  weighs  0.08952289   gram    (one 
crith);  specific  gravity,  0.06949   (Crafts,   C.  r.,  1888,    106,   1662).     It  is  used   for 
filling  balloons;  also  illuminating  gas,  containing  about  50  per  cent  of  hydrogen, 
is  frequently  used  because  it  is  much  cheaper.     It  is  a  non-poisonous  gas,  but 
causes  death  by  exclusion  of  air.     It  has  been  liquified  to  a  colorless  trans- 
parent liquid  by  cooling  to  — 220°  under  great  pressure  and  then  allowing  to 
expand  rapidly  (Olszewski,  C.  r.,  1884,  99,  133;  1885,  101,  2:J8;  Wroblewski,  C.  r., 
1885,  100,  979).     Critical  temperature,  — 234.5°;  critical  pressure,  20  atmospheres; 
tailing  point,  —243.5°    (Olszewski,   Phil.   Mag.,   1895,    (5),   40,   202).     It   diffuses 
through  walls  of  paper,   porcelain,   heated   platinum,   iron,   and  other   metals 
more  than  any  other  gas  (Cailletet,  C.  r.,  1864,  58,  327  and  1057;  1865,  60,  344; 
1868,   66,   847).     It  is   absorbed  by   charcoal   and   by  many   metals,    especially 
palladium;   which,   heated   to    100°    in    an   atmosphere   of    hydrogen   and    then 
cooled  in  that  atmosphere,  absorbs  at-  ordinary  temperatures  982.14  volumes  of 
hydrogen   (Graham,  J.   C.,   1869,  22,   419).     This  occluded  hydrogen   acts  as   a 
strong  reducing  agent,  reducing  FeCl3  to  FeCl2  ,  HgCl2  to  Hg°  ,  etc.     It  is  a 
better  conductor  of  sound  than  air  (Bender,  B.,  1873,  6,  665).     It  conducts  heat 
seven  times  better  than  air  or  480  times  poorer  than  iron  (Stefan,  C.  C.,  1875, 
529).     It  refracts   light  more  powerfully  than   any   other   gas   and   about   six 
times  more  than  air.     It  burns  with  a  non-luminous  flame  and  with  generation 
of  much  heat  (more  than  an  equal  weight  of  any  other  substance  or  mixture 
of    substances).     Hydrogen   forms    two    oxides:    water,    H20  ,    and    hydrogen 
peroxide,  H2O2  (§244). 

2.  Occurrence. — In  volcanic  gases  (Bunsen,  Pogg.,  1851,  83,  197).     In  pockets 
of  certain  Stassfurt  salt  crystals  (Precht,  B.,  1886,  19,  2326).     As  a  product  of 
the  decay  of  organic  material,  both  animal  and  vegetable.     In  combination  as 
water  and  in  innumerable  minerals  (H2O  and  OH)  and  in  organic- compounds. 

3.  Formation. — (a)  By. the  reaction  of  alkali  metals  with  water,     (b]  By 
the  action  of  superheated  steam  upon  heated  metals  or  glowing  coals 
(§226,  4a).     (c)  By  dissolving  aluminum  or  certain  other  metals  in  the 
fixed  alkalis,     (d)  By  the  action  of  many  metals  with  dilute  acids  (seldom 
HN03).     By  heating  potassium  formate  or  oxalate  with  KOH  :   K2C204  + 
2KOH  =  2K2C03  +  H2  (Pictet,  A.  Ch.,  1878,  (5),  13,  216). 

4.  Preparation. — (a)  By  the  action   of   dilute   sulphuric  acid  (one   to 
eight)  on  commercial  or  platinized  zinc  *  (§135,  5a).     The  solution  must 
be  kept  cold  or  traces  of  S02  and  H2S  will  be  evolved.     (&)  By  the  elec- 
trolysis of  acidulated  water. 

5.  Solubilities. — Water    at   ordinary   temperature    dissolves   nearly   two   per 
cent  (volume)  of  hydrogen.     Charcoal  dissolves  or  absorbs  fully  ten  times  its 
volume  of  the  gas  (1). 

6.  Reactions. — Hydrogen   gas  is   a   very  indifferent   body   at   ordinary   tem- 
perature,  combining  with   no  other   element   except    as  it   :s   occluded   or   ab- 
sorbed by  palladium,  platinum,  iron,  nickel,  etc.;  and  in  the  sunlight  combines 
with  chlorine  and  bromine.     "  Nascent  hydrogen  "   (hydrogen  at  the  moment 
of  its  generation),  however,  is  a  powerful  reducing  agent,  and  under  proper 

*  For  the  rapid  generation  of  hydrogen  the  zinc  should  be  granulated  by  pouring  the  molten 
metal  into  cold  water.  Chemically  pure  zinc  is  very  slowly  attacked  by  dilute  sulphuric  acid ; 
but  the  commercial  zinc  frequently  contains  sufficient  impurities  to  insure  a  rapid  generation 
of  hydrogen  when  treated  with  the  dilute  acid.  By  tho  addition  to  the  granulated  zinc,  in  a  tub 
of  water,  of  a  few  cubic  centimetres  of  a  dilute  solution  of  platinum  chloride  or  copper  sul- 
phate, the  zinc  is  made  readily  soluble  in  dilute  sulphuric  acid  and  a  uniform  and  rapid  gen- 
eration of  hydrogen  can  be  obtained. 


§219,  9.  HYDROGEN.  251 

conditions  combines  with  0  ,  S  ,  Se  ,  Te  ,  Cl ,  Br  ,  I ,  N  ,  P  ,  As  ,  Sb  and  Si 
with  comparative  readiness.  The  reduction  of  salts  by  nascent  hydrogen  in 
acid  or  alkaline  solution  will  not  be  discussed  here.  See  under  the  respective 
elements.  It  should  be  noted,  however,  that  "  nascent  hydrogen  "  generated 
by  different  methods  does  not  possess  the  same  reducing  properties.  Sodium 
amalgam  with  acids  does  not  give  hydrogen  capable  of  reducing  silver  halides; 
the  reduction  is  rapid  when  zinc  and  acids  are  used.  Neither  electrolytic 
hydrogen  nor  that  from  sodium  amalgam  and  acids  reduces  chlorates;  while 
zinc  and  acids  reduce  rapidly  to  chlorides.  Hydrogen  generated  by  KOH  and 
Al  does  not  reduce  AsV;  that  formed  by  zinc  and  acids  gives  AsH3  .  Sbv 
with  sodium  amalgam  and  acids  gives  Sb°;  with  zinc  and  acids,  SbH3  (Cha- 
brier,  C.  r.,  1872,  75,  484;  Tommasi,  BL,  1882,  (2),  38,  148). 

Hydrogen  occluded  in  metals  as  Pd  ,  Pt ,  etc.,  is  even  more  active  than 
"  nascent  hydrogen  ";  often  causing  combination  with  explosive  violence 
(Berthelot,  A.  Ch.,  1883,  (5),  30,  719;  Berliner,  W.  A.,  1888,  35,  781).  Hydrogen 
absorbed  by  palladium  precipitates  Ag" ,  Au  ,  Pt ,  Pd  ,  Cu  and  Hg  from  their 
solutions;  permanganates  acidified  are  reduced  to  Mn";  Fe'"  to  Fe";  Crvi  to 
Cr'";  KC103  to  KC10;  CH3CO2H  to  CH3CHO  and  C2H5OH;  and  O6H31TC>2  to 
C6H5NH2  .  The  reactions  are  quantitative.  Salts  of  Pb  ,  Bi  ,  Cd ,  As  ,  Sb  ,  W, 
Mo  ,  Zn  ,  Co  ,  Ni  ,  Al ,  Ce  ,  U  ,  Bb  ,  Cs  ,  K  ,  Na  ,  Ba  ,  Sr  and  Ca  are  not  reduced 
(Schwarzenbach  and  Kritschewsky,  Z.,  1886,  25,  374).  In  the  presence  of 
platinum  black  hydrogen  reduces  very  much  as  described  above;  also  K3Fe(CN)0 
becomes  K4Fe(CN)fi;  dilute  HNO3  becomes  NH4NO2  ,  concentrated  HNO3  be- 
comes HN02;  Cl ,  Br  and  I  combine  with  the  hydrogen  in  the  dark;  KC103 
and  KC10  are  reduced  to  chlorides,  KC104  is  not  reduced;  H2S04  ,  concen- 
trated, is  reduced  to  H2SO3  (Cooke,  C.  N.,  1888,  58,  103). 

Free  hydrogen  very  slowly  acts  upon  a  neutral  solution  of  silver  nitrate, 
precipitating  traces  of  silver;  and  in  concentrated  solution  with  formation  of 
Ag-NO, ;  hindered  by  HN03  or  KNO3  .  Solutions  of  Au  ,  Pt  and  Cu  are  also 
acted  upon  (Rusself,  J.  C.,  1874,  27,  3;  Leeds,  B.,  18715,  9,  1456;  Eeichardt,  Arch. 
Pharm.,  1883,  221,  585;  Toleck  and  Thuemmel,  B.,  1883,  16,  2435;  Senderens,  BL, 
1897,  (3),  15,  991).  KMnO4  in  acid,  neutral,  or  alkaline  solution  slowly 
oxidizes  hydrogen.  It  is  not  at  all  oxidized  by  nitrohj^drochloric  acid,  in 
diffused  daylight,  CrO3  ,  at  ordinary  temperature,  FeCl3  ,  K3Fe(CN)6  ,  HNO3  , 
sp.  gr.  1.42,  or  H2SO4  ,  sp.  gr.  1.84  (Wanklyn  and  Cooper,  Phil.  Mag.,  1890,  (5), 
30,  431).  In  some  cases,  when  hydrogen  under  ordinary  conditions  is  without 
action,  if  subjected  to  great  pressure  a  reducing  action  takes  place;  e.g., 
hydrogen  at  100  atmospheres  pressure  precipitates  Hg°  from  HgCL  (Loewen- 
thal,  J.  pr.,  1860,  79,  480). 

7.  Ignition. — Chlorine  and  bromine  combine  with  hydrogen  directly  in 
the  sunlight,  but  heat  is  required  to  effect  its  combination  with  iodine, 
fluorine,  and  oxygen. 

All  oxides,  hydroxides,  nitrates,  carbonates,  oxalates,  and  organic  salt ; 
of  the  following  elements  are  reduced  to  the  metallic  or  elemental  state  by 
ignition  in  hydrogen  gas :  Pb  ,  Ag ,  Hg ,  Sn  ,  Sb  ,  As ,  Bi  ,  Cu  ,  Cd  ,  Pd  , 
Mo ,  Ru ,  Os ,  Rh ,  Ir ,  Te  ,  Se ,  W ,  Fe ,  Cr ,  Co ,  Ni ,  Zn ,  Tl ,  Nb ,  In  ,  V  . 

Compounds  of  aluminum,  manganese,  and  of  the  fifth  and  sixth  group 
metals  have  not  been  reduced  by  hydrogen. 

8.  Detection. — (a)  Method   of  formation  if  known.     (&)  Its  explosive 
union  with  oxygen  when  the  mixture  with  air  is  ignited,     (c)  Absorption 

.  by  palladium  sponge,  (d)  Explosive  union  with  chlorine  in  the  sunlight 
to  form  HC1 .  (e)  Separated  from  most  other  gases  by  its  non-absorption 
by  the  chemical  reagents  used  in  gas  analysis. 

9.  Estimation. — By  volume   measurement,   almost  never  by   weight,    except 
vrhen  determined  in  its  compounds  by  combustion  to  H.O  . 


BORON— BORIC  ACID.  §220. 

§220.  Boron.     B  =  11.0  .     Valence  three  (§2). 

Boron  does  not  occur  free  in  nature.  It  is  found  chiefly  as  borax,  Na2B407  , 
and  as  boric  acid,  H3B03  ,  in  volcanic  districts.  Two  varieties  of  the  element 
have  been  prepared,  amorphous  and  crystalline.  The  former  is  changed  to  the 
latter  by  heating-  to  a  white  heat  in  presence  of  Al  and  C  (Woehler  and  Claire- 
Deville,  A.,  1867,  141,  268).  Elemental  boron  is  prepared  (a)  by  electrolysis; 
(&)  by  fusing  B203  with  Al ,  Na  or  Mg-;  (c)  by  igniting  BC13  with  hydrogen; 
(d)  by  fusing  borax  with  red  phosphorus.  Specific  gravity  of  the  crystalline, 
2.53  to  2.68  (Hampe,  A.,  1876,  183,  75);  of  the  amorphous,  2.45.  Amorphous 
boron  is  a  greenish-brown,  opaque  powder,  odorless,  tasteless,  insoluble  in 
water,  alcohol  or  ether.  It  is  a  non-conductor  of  electricity.  Heated  in  air  or 
oxygen  it  burns  wTith  incandescence.  In  air  it  forms  B.,O3  and  BN  .  It  is 
oxidized  by  molten  KOH  or  PbCr04  ,  with  incandescence.  It  is  dissolved  by 
concentrated  HN03  or  H2S04  ,  forming  boric  acid.  At  a  red  heat  it  decom- 
poses steam.  When  heated  it  combines  directly  with  S  ,  Cl ,  Br  ,  N  and  many 
metals.  It  forms  BC13  with  chlorine,  not  BC15  .  Fused  with  P20r>  it  forms 
B203  and  P;  with  KOH  ,  K,BO,  and  H;  with  K2C03  ,  K3B03  and  C  .  Boron 
forms  but  one  oxide,  B,03  ,  boric  anhydride.  Three  hydroxides  are  known: 
2H3BO3  =:  B2O3.3H2O  ,  orthoboric  acid;  2HBO2  =  B2O3.H,0  ,  metaboric  acid;  and 
H3B4OT  =  2B2O8,H20  ,  pyroboric  acid. 


§221.  Boric  acid.     H3B03  =  62.024  . 

H'3B'"0-"3  ,  H  —  0  —  B  ~  °  ~  £ 

1.  Properties. — Boron   trioxide,  B,03  ,   boric  anhydride,   is   a  brittle   vitreous 
mass;  sp.  or.  at  12°,  1.8476  (Ditte,  A.  Ch.,  1878,  (5),  13,  67).     Melting  point,  577° 
(Carnelley,  J.  C.,  1878,  33,  278).     It  is  volatile  at  a  very  high  heat  (Ebelemen, 
A.  Ch.,  1848,   (3),  22,  211).     It  has  a  slightly  bitter  taste,  is  hygroscopic,  and 
shows  a  marked  rise  in  temperature  on  solution  in  water   (Ditte,  C.  r.,  1877, 
85,  1069).     In  some  respects  boron  trioxide  deports  itself  as  a  weak  base.     It 
forms  a  sulphide,  B,S3  ,  decomposed  by  water    (Woehler  and  Deville,  A.   Ch., 
1858,  (3),  52,  90);  a  sulphate,  B(HS04)3    (D'Arcey,  J.  C.,  1889,  55,  155);  and  a 
phosphate,  BPO4   (Meyer,  B.,  1889,  22,  2919).     It  combines  with  water  in  three 
proportions,  forming  the  ortho,  meta  and  pyroboric  acids.     Orthoboric  acid  is 
a  weak  acid,  its  solutions  reddening  litmus;  at  12°  it  has  a  specific  gravity  of 
1.5172  (Ditte,  I.e.);  melts  at  184°  to  186°   (Carnelley,  I.  c.).     Soluble  in  25  parts 
water  at  20°,  and  in  3.4  parts  at  102°  (Ditte,  I.e.).     It  is  volatile  in  steam  and 
in  alcohol  vapor.     The   evaporation  of  the  water  of  combination  of  the  acid 
carries  with  it  from  ten  to  fifteen  per  cent  of  the  acid. 

2.  Occurrence. — Widely  distributed,  but  usually  in  very  small  quantities.     In 
the  rock  salt  deposits  at  Stassfurt,  Germany,  as  boracite,  Mg7B]flOsoCl2    (62.5 
per  cent  B2OS).     In  the  volcanic  regions  of  Tuscany  and  the  Liparic  Islands  as 
steam  saturated  with  boric  acid. 

3.  Formation. — The  anhydride  is  formed  by  burning  the  metal  in  air 
or  oxygen,  or  by  heating  the  acids.     Orthoboric  acid,  H3B03 ,  is  formed 
by  dissolving  the  oxide  in  water;  the  meta  acid,  HB02 ,  H  —  0  —  B  =  0 , 
by  heating  the  ortho  acid  a  little  above  100°  (Bloxam,  /.  C.,  1860,  12, 
177);  the  pyroboric  acid,  tetraboric  acid,  H2B407 ,  by  heating  the  ortho 
or  meta  acid  for  some  time  at  160°  in  a  current  of  dry  air  (Merz,  J.  pr., 
1866,  99,  179). 

4.  Preparation. — (a)  By  evaporation  of  the  water  from  the  lagoons  of 
Tuscany,    which    are    saturated    with    boric    acid,    and    recrystallization 


§221,  7.  BORIC  ACID.  253 

from  water,  (b)  The  boronatrocalcite,  Ca,B6On.Na2B407  +  18H20  (45.6 
per  cent  B203),  found  in  Nevada,  is  evaporated  in  lead  pans  with  H2S04  to  a 
stiff  paste;  and  then  treated  with  superheated  steam  in  iron  cylinders 
heated  to  redness.  The  acid  passes  over  with  the  steam  and  is  collected 
in  lead  lined  chambers  (Gutzkow,  Z.,  1874,,  13,  457).  (c)  Commercial 
borax,  Na2B407.10H20 ,  is  dissolved  in  hot  water,  twelve  parts,  and  acidi- 
fied with  hydrochloric  acid.  Upon  cooling,  the  boric  acid,  H3B03 ,  is  ob- 
tained in  small  scales,  which  are  purified  by  recrystallizatioii  from  hot 
water. 

5.  Solubilities. — More  soluble  in  hydrochloric  acid  solution  or  in  alcohol 
than  in  water  (1).     The  alcoholic  solution  burns  with  a  beautiful  green 
flame.     Quite  soluble  in  glycerine  and  in  most  alcohols  and  hydrocarbons, 
only  sparingly  in  ether.     The  borates  are  insoluble  in  alcohol;  those  of 
the  alkalis  are  soluble  in  water  to  an  alkaline  solution.     Borates  of  the 
other  metals  are  insoluble  in  water  (no  borate  is  entirely  insoluble  in 
water) ;  but  are  usually  rendered  soluble  by  the  addition  of  boric  acid. 

6.  Reactions. — Silver  nitrate  forms,  in  solutions  of  acid  borates,  a  white 
precipitate  of  silver  meiaborate,  AgB02 ,  but  normal  borates   form  in  part 
silver  oxide,  brown.     Lead  acetate  gives  a  white  precipitate  of  lead  meta- 
lorate,  Pb(B02)2  ;  calcium  chloride,  in  solutions  not  very  dilute,  a  white 
precipitate  of  calcium  metaborate;    and  barium  chloride,  in  solutions  not 
dilute,    a    white    precipitate    of    barium    metaborate,    Ba(B02)2.      With 
aluminum  salts,  the  precipitate  is  aluminum  hydroxide. 

Borates  are  transposed  with  formation  of  boric  acid,  by  all  ordinary 
acids — in  some  conditions  even  by  carbonic  acid. 

The  liberated  boric  acid  is  dissolved  by  alcohol,  and  if  the  alcohol  solu- 
tion be  set  on  fire,  it  burns  with  a  green  flame. 

A  solution  of  a  borate,  acidulated  with  hydrochloric  acid  to  a  barely 
perceptible  acid  reaction,  imparts  to  a  slip  of  turmeric  paper  half  wet  with 
it,  a  dark-red  color,  which  on  drying  intensifies  to  a  characteristic  red  color 
which  turns  dark  green  when  moistened  with  a  drop  of  alkali. 

7.  Ignition. — Boric  acid  is  displaced  from  its  salts  by  nearly  all  acids 
including  C02  ;  but  being  non-volatile  except  at  a  very  high  heat,  it  dis- 
places most  other  acids  upon  ignition. 

By  heating  a  mixture  of  borax,  acid  sulphate  of  potassium,  and  a  fluo- 
ride, fused  to  a  bead  on  the  loop  of  platinum  wire,  in  the  clear  flame  of 
the  Bunsen  gas-lamp,  an  evanescent  yellowish-green  color  is  imparted  to 
the  flame. 

Borates  fused  in  the  inner  How-pipe  flame  with  potassium  acid  sulphate 
give  the  green  color  to  the  outer  flame. 

If  a  crystal  of  boric  acid,  or  a  solid  residue  of  borate  previously  treated 
with  sulphuric  acid,  on  a  porcelain  surface,  is  played  upon  by  the  flame  of 
Bunsen's  Burner,  the  green  flame  of  boron  is  obtained. 


254  CARBON.  §221,  8. 

If  a  powdered  borate  (previously  calcined),  is  moistened  with  sulphuric 
acid  and  heated  on  platinum  wire  to  expel  the  acid,  then  moistened  with 
glycerine  and  burned,  the  green  flame  appears  with  great  distinctness. 
The  glycerine  is  only  ignited,  then  allowed  to  burn  by  itself.  Barium 
does  not  interfere  (being  held  as  sulphate,  non-volatile);  copper  should  be 
previously  removed  in  the  wet  way.  The  glycerine  flame  gives  the  spec- 
trum. But  in  all  flame  tests,  the  boric  acid  must  be  liberated. 

Borates  (fused  on  platinum  wire  with  sodium  carbonate)  give  a  char- 
acteristic spectrum  of  four  lines,  equidistant  from  each  other,  and  extend- 
ing from  Ba  Y  in  the  green  to  Sr  d  in  the  blue. 

Borax,  Na2B407 ,  when  ignited  (as  on  a  loop  of  platinum  wire  to  form 
the  borax  bead)  with  many  metallic  compounds,  forms  a  colored  glass, 
used  in  the  detection  of  certain  metals  (§132,  7).  The  fused  borax  forms 
a  solid  brittle  mass,  borax  glass,  used  in  assaying  and  in  soldering  because 
of  its  power  of  combination  with  metallic  oxides. 

8.  Detection. — By  conversion  into  the  acid,  if  present  as  a  salt;  solution 
in  alcohol  or  glycerine  and  burning  with  the  formation  of  the  green  flame 
(very  delicate,  but  copper  salts  should  be  removed  by  H2S  and  barium  salts- 
should  be  removed  or  converted  into  the  sulphate).     Also  by  the  red  color 
imparted  to  a  strip  of  turmeric  paper. 

9.  Estimation. — Boron  compounds   cannot  be   completely  precipitated   from 
solution  by  any  known  reagents,  hence  most  of  the  methods  of  quantitative 
determination  are  indirect.     By  adding  a  known  quantity  of  Na2C03  ,  fusing 
and   weighing;   then   after   determining  the   CO2    subtracting   its   weight   and 
that  of  the  Na,O  present   (calculated  from  Na,C03  first  added).     The  differ- 
ence is  the  weight  of  B2O3  present.     See  also  Will  (Arch.  Plianu.,  1887,  225,  1101). 
In  the  presence  of  glycerine  or  mannitol,  boric  acid  may  be  accurately  titrated 
with  sodium  hydroxide,  using  phenolphthalein  as  an  indicator:   B2O3  +  2NaOH  = 

:  NaBO2  +  H2O  .  Sodium  carbonate  must  be  absent  or  we  get:  2B2O3  +  Na2CO.-i  = 
Na^B.Oy  +  COo  (Honig  and  Spitz,  Z.  angew.,  1896,  549;  Joergensen,  Z.  angew., 
1897,  5). 


§222.  Carbon.     C  =  12.0  .     Usual  valence  four  (§15). 

1.  Properties. — Carbon  exists  in  three  allotropic  forms:  two  crystalline, 
diamond  and  graphite,  and  amorphous  as  charcoal,  coke,  etc.  Specific  gravity, 
diamond  at  4°,  3.51835  (Baumhauer,  J.,  1873,  237);  graphite,  Ceylon,  2.2o'to  2.26 
(Brodie,  A.,  1860,  114,  6);  wood  charcoal,  1.57;  gas  coke,  1.88.  Very  small 
specimens  only,  of  diamonds  have  been  artificially  prepared,  by  saturating  iron 
with  carbon  at  3000°.  At  this  temperature  graphite  is  formed  and  upon  cool- 
ing under  pressure  the  crystalline  diamond  form  is  obtained.  This  cooling 
under  pressure  is  obtained  by  pouring  the  carbon  saturated  iron  into  a  soft 
iron  bomb,  which  is  cooled  by  water  (Moisson,  C.  r.,  1893,  116,  218).  Diamond 
is  the  hardest  substance  known.  It  is  very  strongly  refractive  towards  light 
(Becquerel,  A.  C7».,  1S77,  (5),  12,  5).  Fluorescence  and  phosphorescence  of 
diamonds,  see  Kunz  (C.  (7.,  1891,  ii,  562).  Ignition  in  an  atmosphere  of  hydro- 
gen does  not  effect  a  change;  in  air  or  oxygen  it  burns  to  CO.,  . 

Graphite  is  a  hard,  gray,  metal-like,  opaque  solid,  a  good  conductor  of 
electricity  and  a  fairly  good  conductor  of  heat.  It  burns  with  difficulty.  It 


§222,  0.  CARBON.  255 

is  used  in  lead  pencils,  in  black  lead  (plumbago)  crucibles,  as  a  lubricant  for 
heavy  machinery,  in  battery  plates,  lor  The  arc  lig-ht  carbon  pencils,  etc. 

Amorphous  carbon  is  black,  lighter  than  diamond  or  graphite.  It  is  in  use 
as  coal,  coke,  charcoal,  animal  charcoal,  etc.;  all  impure  forms.  Lamp-black 
is  also  amorphous  carbon  made  from  burning  resin,  fat,  wax,  coal  gas,  etc., 
with  limited  supply  of  air.  It  is  used  as  a  pigment  in  paints,  in  stove-black- 
ing, shoe-blacking,  printers'  ink,  etc.  Charcoal,  preferably  animal  charcoal,  is 
used  for  decoloring  organic  solutions.  Charcoal  absorbs  many  gases,  hence  is 
valuable  as  a  disinfectant. 

Carbon  forms  two  oxides:  carbon  monoxide,  CO  ,  and  carbon  dioxide,  C02  . 

2.  Occurrence. — Diamonds  seem  first  to  have  been  found  in  India,  especially 
in  the  Golconda  pits,  where,  as  early  as  1622,  30,000  laborers  are  said  to  have 
been  employed  (Walker,  /.,  1884,  774).  Also  found  in  other  parts  of  Asia,  in 
South  Africa,  in  Brazil,  etc.  (Winklehner,  C.  C.,  1888,  192;  Damour,  J.,  1883,  774; 
Gorceix,  J.,  1881,  345;  Smit,  J.,  isso,  1400).  Graphite  is  found  in  Ceylon  (Wal- 
ther,  C.  C.,  1890,  ii,  20);  in  California  (C.  N.,  1868,  17,  209);  in  Canada  (Dawson, 
Am.  8.,  1870,  (2),  50,  130);  in  New  York  State;  in  New  Zealand  (Mac  Ivor, 
(\  N.,  1887,  55,  125);  in  Russia,  Germany,  Greenland,  etc.  Pure  amorphous 
r—^on  occurs  in  nature  a-;  a  chief  product  in  the  decomposition  of  organic 
material,  air  being  excluded.  Anthracite  coal  is  relatively  pure  amorphous 
carbon. 

3.  Formation. —  Graphite  remains  as  a  residue  when  pig  iron  is  dis- 
solved in  acids.     It  forms  by  reducing  CO  with  Fe304  at  400°.     Amor- 
phous carbon  is  formed  by  passing  CC14  over  Na  in  a  tube  heated  to  red- 
ness (Porcher,  C.  A7.,  1881,  44,  203). 

4.  Preparation. — Pure  graphite  is  prepared  by  heating  the  commercial 
graphite  on  a  water  bath  with  KC103  and  H2S04  and  repeatedly  washing. 
If  it  contains  Si02  it  should  also  be  treated  with  NaF  and  H2S04 .     Amor- 
phous carbon  is  prepared  by  heating  wood,  coal,  or  almost  any  organic- 
matter  to  a  very  high  temperature  in  absence  of  air,  but  when  so  prepared 
it  is  never  pure.     Amorphous  carbon  is  prepared  approximately  pure  by 
heating  pure  cane  sugar  in  a  closed  platinum  crucible;  then  boiling  in 
succession  with  HC1 ,  KOH ,  and  H20  ;  then  igniting  to  redness  in  an 
atmosphere  of  chlorine,  cooling  in  the  same  atmosphere.     A  very  pure 
form  of  graphite  known  as  Acheson  graphite  is  made  in  large  quantities  at 
Niagara  Falls  by  passing  a  strong  electric  current  through  coke  compressed 
into  the  desired  form.     Silicon,  aluminum  and  other  impurities  are  dis- 
tilled out  at  the  high  temperature  employed. 

5.  Solubilities. — Insoluble  in  water  or  acids       Soluble  in  many  molten 
metals  with   partial  combination  to  form  carbides.     When  the  metal  is 
dissolved  in   acids  the  combined  carbon   passes  off  as  hydrocarbons,  the 
excess  remaining  as  graphite. 

6.  Reactions. — Xot  attacked  by  acids  or  alkalis.     It  slowly  oxidizes  to 
C02  when  heated  with  concentrated  H2S04   and  K2Cr20T .     Upon  gently 
warming   graphite   with  KC103   and   HN03 ,   graphitic   acid,   C^H^,   is 
said  to  be  formed  (Stingl,  B.,  1873,  6,  391).     The  important  reactions  of 
carbon  require  the  aid  of  high  heat  and  are  described  in  the  next  para- 
graph. 


ACETIC  ACID. 


§22$,  ?. 


7.  Ignition.  —  Unchanged  by  ignition  in  absence  of  air.    When  strongly 
ignited  in  air  or  oxygen  it  slowly  burns  to  C02  .    If  the  carbon  and  oxygen 
has    been    previously    very    thoroughly    dried    the    action    is    very    slow, 
especially  with  graphite.     By  fusion  with  KN03  or  KC103  carbon  is  oxid- 
ized to  C02  .    With  vapors  of  sulphur,  carbon  disulphide  is  formed  ;  i.  e.,  by 
passing  sulphur  vapors  over  hot  coals  in  an  electrically  heated  furnace.    In  an 
u'mosphere  of  hydrogen  with  the  electric  spark,  acetylene,  C2H2  ,  is  formed. 
i-y   igniting  in  an  atmosphere  of  carbon  dioxide.  C02  ,  the  whole  of  the 
carhon  becomes  carbon  monoxide:  C  +  C02  =  2CO. 

By  simple  ignition  with  carbon,  all  oxides  of  the  elements  in  the  follow- 
ing list  are  reduced  to  the  elemental  state  (a);  and  if  sodium  carbonate  is 
added,  all  of  the  salts  of  the  same  are  likewise  reduced  (b).  Cu  ,  Bi  ,  Cd  , 
Pb,  Ag,  Hg,  As,  Sb,  Sn,  Pd,  Mo,  Ru  ,  Os,  Rh,  Ir,  Te,  Se,  w',  K, 
Na  ,  Rb  ,  Cr  ,  Fe  ,  Mn  ,  Co  ,  Ni  ,  Zn  ,  Ti  ,  Tl  . 

(a)     Pb304  +  2C  =  3Pb  +  2CO2 

(6)     2PbCL  +  2Na,CO3  +  C  =  2Pb  +  4NaCl  +  3C02 

(c)  CuO  +  C  (excess)  =  Cu  +  CO 

(d)  C  -f-  2CuO  (excess)  =  2Cu  +  CO. 

With  excess  of  carbon  CO  is  formed  (c).  With  excess  of  the  oxide  C02  is 
formed  (d).  In  the  reduction  of  iron  ore,  the  process  is  conducted  so  as 
to  give  some  CO  and  some  C02  .  To  obtain  some  metals  in  the  free  state 
(such  as  K  and  Na),  special  methods  are  adopted  to  exclude  the  air,  and 
to  produce  the  high  temperature  needed. 

All  compounds  of  sulphur  when  ignited  with  carbon  are  reduced  to  a 
sulphide:  BaS04  +  2C  =  BaS  +  2CO,  . 

8.  Detection.  —  By  its  appearance;  failure  to  react  with  general  reagents; 
and  by  its  combustion  to  C02  with  oxygen  (air),  or  with  K2Cr207  and  con- 
centrated H2S04  (Fritsche,  A.,  1896,  294,  79),  then  by  identification  with 
Ca(OH)2  (§228,  6). 

9.  Estimation.  —  By  combustion  to  CO,  and  weighing1  after  absorption  in  KOH 
solution.  See  works  on  ultimate  organic  analysis. 


§223.  Acetic  acid.     HC2H302  =  60.032  . 

H      0 

I        II 
H'4(C2)+"'-"'0-"2 ,  H  — C  — C  — 0  — H  =  CH3C02H. 


1.  Properties. — Pure  acetic  acid  is  a  colorless,  crystalline,  hygroscopic  solid, 
melting  at  16.5°  and  boiling  at  118°.  Its  specific  gravity  at  0°  is  1.080.  It  has 
a  sharp,  sour  taste,  an  irritating  burning  ell'ect  on  the  skin,  and  a  very  peiie- 


§223,  6.  ACETIC   ACID.  257 

trating  odor.  It  burns  when  heated  nearly  to  the  boiling  point.  Vinegar  contains 
four  to  five  per  cent  of  acetic  acid.  The  U.  S.  P.  reagent  contains  36  per  cent  of 
acetic  acid,  and  has  a  specific  gravity  of  1.0481  at  15°.  It  vaporizes  from  its  con- 
centrated solutions  at  ordinary  temperature,  having  the  characteristic  odor  of 
vinegar.  It  is  a  monobasic  acid,  as  the  three  remaining  hydrogen  atoms  (linked 
to  carbon)  cannot  be  replaced  by  metals. 

2.  Occurrence. — It  occurs  in  nature  in  combination  with  alcohols  in  the  essen- 
tial oils  of  many  plants. 

3.  Formation.— (a)  During  the  decay  of  many  organic  compounds.     (&) 
By  gently  heating  sodium  methylate,  NaOCH3 ,  in  a  current  of  carbon 
monoxide:    NaOCH3    +    CO    =    CH3C02Na   (NaC2H302).     (c)  By   boiling 
methyl  cyanide  with  acids  or  alkalis :   CH3CN  +  HC1  +  2H20  =  HC2H302 
+   NH4C1.     (d)  By  the  oxidation  of  alcohol:    3C2HG0   +   2K2Cr207   + 
8H2S04  ==  2K2S04  +  2Cr2(S04)3  +  3HC2H30,  +  11H20  . 

4.  Preparation. — (a)  By  the  dry  distillation  of  wood.     (&)  By  the  fer- 
nlentation  of  cider,  beer,  wine,  molasses,   etc.     (c)  Pure  acetic  acid  is 
prepared  by  distilling  anhydrous  sodium  acetate  with  concentrated  sul- 
phuric acid.     The  distillate  solidifies  upon  cooling  and  is  termed  glacial 
acetic  acid. 

5.  Solubilities. — Miscible  in  all  proportions  in  water  and  alcohol.     The 
salts  of  acetic  acid,  acetates,  are  all  soluble  in  water,  silver  and  mercurous 
acetates  sparingly  soluble.     One  part  of  silver  acetate  requires  115  parts 
of  water  at  10°  for  its  solution;  one  part  of  mercurous  acetate  requires 
•  loo  parts  of  water.     Certain  basic  acetates,  as  Fe'" ,  Al ,  etc.,  are  insoluble 
in  water.     Very  many  of  the  acetates  are  soluble  in  alcohol.     Ammonium 
acetate    dissolves   several   insoluble   sulphates   such    as   calcium    and    lead 
sulphates. 

6.  Reactions. — The    stronger    mineral    acids    transpose    the    acetates, 
forming   acetic   acid.     Anhydrous   acetates   with   concentrated   sulphuric 
acid  give  pure  acetic  acid  (4),  but  if  the  sulphuric  acid  be  in  excess  and 
heat  be  applied  the  mixture  blackens  with  separation  of  carbon;  and,  at 
higher  temperatures,  C02  and  S02  are  evolved. 

Solution  of  ferric  chloride  forms,  with  solutions  of  acetates,  a  red  solu- 
tion containing  ferric  acetate,  Fe(C2H302)3 ,  which  on  boiling  precipitates 
brownish-red,  basic  ferric  acetate.  The  red  solution  is  not  decolored  by 
solution  of  mercuric  chloride  (distinction  from  thiocyanate);  but  is  de- 
colored by  strong  acidulation  with  sulphuric  acid  or  hydrochloric  acid  (dis- 
tinction from  thiocyanate  and  from  meconate).  The  ferric  acetate  is  pre- 
cipitated by  alkali  hydroxides. 

If  acetic  acid  or  an  acetate  be  warmed  with  sulphuric  acid  and  a  little 
alcohol,  the  characteristic  pungent  and  fragrant  odor  of  ethyl  acetate  or 
acetic  ether  is  obtained: 

HC2H,0,  +  C2H5OH  =  H20  +  aH5C2H3O2 

Acetic  acid  does  not  act  as  a  Reducing  Agent  as  readily  as  do  most  of 


258  CITRIC  ACID.  §223,  7. 

the  organic  carbon  compounds.  Acetic  acid  is  stable  toward  mild  oxidiz- 
ing agents  and  is  only  slowly  attacked  by  strong  oxidizing  agents  such  as 
chromic  acid  and  potassium  permanganate;  reduces  auric  chloride  only 
in  alkaline  solution,  and  does  not  reduce  alkaline  copper  solution.  It 
takes  chlorine  into  combination — slowly  in  ordinary  light,  quickly  in  sun- 
light, forming  chloracetic  acids. 

Acetic  acid  is  a  weak  acid,  only  0.4%  being  ionized  in  normal  solu- 
tion and  1.4%  in  tenth  normal  solution.  Glacial  acetic  acid  dissolves 
sulphur  which  crystallizes  out  in  needles. 

7.  Ignition. — By    ignition  alone,    acetates    blacken,    with   evolution    of 
vapor  of  acetone,  C3H60 ,  inflammable  and  of  an  agreeable  odor.     By  pro- 
longed ignition  of  alkali  acetates  in  the  air,  carbonates  are  obtained  free 
from   charcoal.      By   ignition   with    alkali   hydroxides   in    dry   mixtures, 
methane,    marsh-gas,    CH4 ,    is    evolved.      By    ignition    with  alkalis  and 
arsenous  anhydride,  the  poisonous  and  offensive  vapor  of  cacodyl  oxide 
is  obtained.     This  test  should  be  made  under  a  hood  with  great  caution 
and    with    small    quantities.      It    is    a    very    delicate    test    for    acetates: 
4KC2H302  +  As203  =  As,(CH3)40  +  2K2C03  +  2C02. 

Butyric,   propionic  and  valerianic  acids  give  the  same  reaction. 

8.  Detection. — (a)  By  its  odor.     (5)  By  the  formation  of  the  fragrant 
ethyl  acetate  upon  warming  with  sulphuric  acid  and  alcohol.     Too  much 
alcohol  should  not  be  used  in  testing  for  acetic  acid  as  otherwise  ethyl 
ether  is  formed,      (c)    By  the  formation  of  the  red  solution  with  ferric 
chloride   (§126,  6b  and  §152).   ('d)   By  ignition  of  the  dry  acetate  alone 
to  acetone,  CH3COCH3  ;  with  NaOH  to  methane,  CH4  ;  or  with  As,03  to 
cacodyl  oxide,     (e)  As  a  delicate  test  for  formates  or  acetates  it  is  directed 
to  warm  a  solution  of  CuCL  in  NaCl  and  add  a  small  amount  of  the  mate- 
rial under  examination.     Formates  give  a  blackish-gray  deposit;  acetates 
give  a  bright  green  precipitate  not  changed  by  boiling.     Both  precipitates 
are  soluble  in  acetic  acid  (Field,  J.  C.f  1873,  26,  575). 

9.  Estimation.— Other  volatile  acids  are  separated  by  precipitation;  sulphuric 
acid  is  then  added  and  the  acetic  acid  is  distilled  into  water  and  estimated 
by  titration  with  standard  alkali. 


§224.  Citric  acid.     H3C6H507  =  192.064  . 

H2C  —  C02H 

H'3(C6)+10-4H'50-"7  ,  H  —  0  —  C  —  CO  ,H 

I 
HL>C  —  C02H 

Found  in  small  quantities  in  the  juices  of  many  fruits.     The  chief  commercial 
source  is  lemon-juice.     It  is  a  colorless,  crystallizable,  non-volatile  solid;    it  ig 


§225,  1.  TARTARIC  ACID.  259 

soluble  in  0.75  parts  of  cold  water,  in  equal  parts  of  90  per  cent  alcohol,  in  1.3 
parts  of  absolute  alcohol  and  in  50  parts  of  ether.  It  crystallizes  with  one  mole- 
cule of  water  in  rhombic  prisms  which  melt  at  100°. 

The  citrates  of  the  metals  of  the  alkalis  are  freely  soluble  in  water;  those 
of  iron  and  copper  are  moderately  soluble;  those  of  the  alkaline  earth  metals 
insoluble.  There  are  many  soluble  double  citrates  formed  by  action  of  alkali 
citrates  upon  precipitated  citrates,  or  of  alkali  hydroxides  upon  metallic  salts 
in  presence  of  citric  acid.  In  distinction  from  tartrates,  the  solubility  of  the 
potassium  salts,  non-precipitation  of  calcium  salt  in  cold  solution;  and  weaker 
reducing  action,  are  to  be  noted. 

Solution  of  calcium  hydroxide  in  excess  (as  by  dropping  the  solution  tested 
into  the  reagent)  gives  no  precipitate  with  citric  acid  or  citrates  in  the  cold 
(distinction  from  tartaric  acid),  but  on  heating,  the  white  calcium  citrate, 
Ca3(C6H5O7)2 ,  is  precipitated  (not  soluble  in  cold  potassium  hydroxide  solu- 
tion). By  filtering  before  boiling,  the  tartrate  and  citrate  may  be  approxi- 
mately separated.  The  chlorides  of  calcium  and  barium  give  no  precipitate 
in  neutral  solutions  (difference  from  tartaric  acid),  but  if  sodium  hydroxide  is 
added,  calcium  citrate  is  precipitated,  insoluble  in  sodium  hydroxide,  but  readily 
soluble  in  ammonium  chloride.  On  boiling  the  solution  in  ammonium  chloride, 
crystallized  calcium  tartrate  is  precipitated,  which  is  now  insoluble  in  ammo- 
nium chloride.  Calcium  citrate  is  soluble  in  acetic  acid  (distinction  from  oxalates). 

Solution  of  lead  acetate  precipitates  white  lead  citrate,  Pb3(C6H5O7)2 ,  soluble 
in  ammonia.  Silver  nitrate  gives  a  white  precipitate  of  silver  citrate, 
AgsCeH^Or ,  which  does  not  blacken  on  boiling  (distinction  from  tartrate). 
For  action  of  citric  acid  or  citrates  in  hindering  many  of  the  usual  analytical 
reactions,  see  Spiller,  J.  C.,  1858,  10,  110. 

One  part  of  citric  acid  dissolved  in  two  parts  of  water,  and  treated  with  a 
solution  of  one  part  of  potassium  acetate  in  two  parts  of  water,  should  remain 
clear  after  addition  of  an  equal  volume  of  strong  alcohol  (absence  of  oxalic 
acid  and  of  tartaric  acid  and  its  isomers). 

Citric  acid  does  not  act  very  readily  as  a  reducing  agent;  does  not  reduce 
alkaline  copper  solution,  or  silver  solution;  reduces  permanganate  very  slowly. 
Concentrated  nitric  acid  produces  from  it,  acetic  and  oxalic  acids;  and  diges- 
tion with  manganese  dioxide  decomposes  it 'with  formation  of  acetone,  acrylic 
and  acetic  acids.  Concentrated  sulphuric  acid  carbonizes  citric  acid  with  liber- 
ation of  sulphur  dioxide.  Citrates  carbonize  on  ignition,  with  various  empy- 
reuniatic  products,  and  with  final  formation  of  carbonates.  On  heating  citric 
acid,  it  loses  its  water  of  crystallization,  then  fuses  and  decomposes  with  evolution 
of  pungent  fumes  leaving  a  carbonaceous  residue.  By  fused  potassium  hydroxide, 
short  of  ignition,  they  are  decomposed  with  production  of  oxalate  and  acetate. 

§225.  Tartaric  acid.     H2C4H406  =  150.048  . 
H      0 

I        II 
H  —  0  —  C  —  C  —  0  —  H      CH(OH)C<UJ 

I  or  | 

H  —  0  —  C  —  C  —  0  —  H      CH(OH)C00H 

I        II 
H      0 

1.  Properties. — Tartaric  acid  is  a  colorless,  crystalline,  non-volatile  solid; 
soluble  in  0.756  parts  of  water  at  15°,  in  3.4  parts  of  90  per  cent  alcohol,  in  50 
parts  of  ordinary  ether  and  250  parts  of  absolute  ether.  It  exists  in  four  dis- 
tinct modifications:  dextrotartaric  acid,  levotartaric  acid,  racemic  acid,  and 
mesotartaric  acid.  They  differ  from  each  other  in  crystalline  form,  in  solubility. 


260  TARTARIC  ACID.  §225,  2. 

and  especially  in  the  deportment  of  their  solutions  towards  polarized  light. 
Racemic  and  mesotartaric  acids  are  optically  inactive,  but  the  former  may 
he  resolved  into  the  first  two  acids,  optically  active. 

2.  Occurrence. — It  is  found  in  various  fruits.     The  chief  commercial  source 
fs  grape  juice. 

3.  Formation. — By  oxidation  of  dextrose,  cane  sugar,  milk  sugar,  starch,  etc., 
with  HNO3    (Kiliani,   A.,   18SO,   205,    175).     By    action   of   sodium   amalgam   on 
oxalic  ether   in   alcoholic   solution    (Debus,   .{.,    1873,    166,    124).     By    synthesis 
from  succinic  acid  by  formation  first  of  the  dibromsuccinic  acid,  H2C4Br,H204; 
then  substitution  of  the  OH  group  for  the  bromine  by  means  of  water  and 
silver  oxide. 

4.  Preparation. — The  crude  argol  deposited  during  the  fermentation  of  grape 
juice   is   recrystallized,   giving   the    commercial    cream    of   tartar,    KHC4H406  . 
This  in  hot  solution  is  treated  with  powdered  chalk,  and  the  filtrate  from  the 
precipitate   thus   obtained   is    precipitated    with    calcium    chloride.     Both    pre- 
cipitates are  washed  and  decomposed  by  the  necessary  quantity  of  hot  dilute 
sulphuric  acid.     The  tartaric  acid  solution  is  evaporated  to  crystallization  and 
purified  by  recrystallization  (Ficinius,  Arch.  PMrm.,  1879,  215,' 14  and  310). 

5.  Solubilities. — The  Tartrates  of  the  alkali  bases  are  soluble  in  water; 
the  normal  tartrates  being  freely  soluble,  the  acid  tartrates  of  potassium 
and  ammonium  sparingly  soluble.     The  tartrates  of  the  alkaline  earth 
bases  and  of  the  non-alkaline  bases,  are  insoluble  or  sparingly  soluble,  but 
mostly  dissolve  in  solution  of  tartaric  acid.     Most  of  the  tartrates  are 
insoluble  in  alcohol.     There  are  double  tartrates  of  heavy  metals  with 
alkali  metals,  which  dissolve  in  water.     Tartar-emetic  is  potassium  anti- 
mony! tartrate,  KSbOC4H406 . 

Hydrochloric,  nitric,  and  sulphuric  acids  transpose  the  tartrates 
(whether  forming  solutions  o>  not).  Most  of  the  tartrates  are  also  dis- 
solved (and,  if  already  dissolved,  are  not  precipitated)  by  the  alkali  hy- 
droxides, owing  to  the  formation  of  soluble  double  tartrates. 

The  freshly  precipitated  oxides,  hydroxides,  and  carbonates  of  the  fol- 
lowing metals  are  soluble  in  a  solution  of  potassium-sodium  tartrate, 
Rochelle  salt:  Sb  ,  Snlv,  Bi ,  Cu ,  Fe  ,  Al ,  Cr ,  Co  ,  Ni ,  Mn  ,  and  Zn  ;  Ba  , 
Sr,  Ca,  and  Mg  to  quite  an  extent.  CdC03  is  not  dissolved  (Warren, 
C.  N.,  1888,  57,  223). 

6.  Reactions. — Solution  of  calcium  hydroxide,  added  to  alkaline  reac- 
tion, precipitates  from  cold  solution  of  tartaric  acid,  or  of  soluble  tartrates, 
calcium  tartrate,  white,   CaC4H406 .     Solution   of   calcium   chloride   with 
neutral  tartrates  gives  the  same  precipitate.     Solution  of  calcium  sulphate 
forms  a  precipitate  but  slowly,  or  not  at  all  (distinction  from  racemic  acid). 
The  precipitate  of  calcium  tartrate  is  soluble  in  cold  solution  of  potassium 
hydroxide,  precipitated  gelatinous  on  boiling,  and  again  made  soluble  on 
cooling  (distinctions  from  citrate),  and  dissolves  in  acetic  acid  (distinction 
from  oxalate). 

Tartaric  acid  prevents  the  precipitation  by  fixed  alkalis  of  solutions  of 
the  &alts  of  the  following  metals :  Al ,  Bi ,  Co  ,  Ni ,  Cr ,  Cu ,  Fe  ,  Pb ,  Pt , 
and  Zn  (Grothe,  /.  pr.,  1864,  92,  175). 


§225,  9.  TARTARIC  ACID.  261 

Silver  nitrate  precipitates,  from  solutions  of  normal  tartrates,  silver 
tartrate,  Ag2C4H400 ,  white,  becoming  black  when  boiled.  If  the  precipi- 
tate is  filtered,  washed,  dissolved  from  the  filter  by  dilute  ammonium 
hydroxide  into  a  clean  test-tub**,  left  for  a  quarter  of  an  hour  on  the 
water-bath,  the  silver  is  reduced  as  a  mirror  coating  on  the  glass  (§59,  lOb), 
distinction  from  citric  acid.  Free  tartaric  acid  does  not  reduce  silver 
salts.  Permanganate  is  reduced  quickly  by  alkaline  solution  of  tartrates 
(distinction  from  citrates),  precipitating  manganese  dioxide,  brown.  Free 
tartaric  acid  acts  but  slowly  on  the  permanganate.  Alkaline  copper  tar- 
trate,  Fehling's  solution  (§77,  6b),  resists  reduction  in  boiling  solution. 
Chromates  are  reduced  by  tartaric  acid,  the  solution  turning  green.  The 
oxidized  products,  both  with  permanganate  and  chromate,  are  formic 
acid,  carbonic  anhydride,  and  water. 

7.  Ignition. — On  ignition,  tartaric  acid  or  tartrates  evolve  the  odor  of 
"burnt  sugar,  separating  carbon,  and  becoming  finally  converted  to  carbon- 
ates.— Strong  sulphuric  acid  also  blackens  tartrates,  on  warming.    Melted 
potassium  hydroxide,  below  ignition,  produces  acetate  and  oxalate.     The 
fixed  alkali  tartrates  ignited  in  absence  of  air  give  an  alkali  carbonate  and 
finely  divided  carbon.     The  mixture  serves  as  an  admirable  flux  for  the 
reduction  tests  for  arsenic  (§69,  7). 

8.  Detection. — (a)  By  the  odor  of  burnt  sugar  when  ignited.     (&)  By 
the  deportment  of  the  calcium  salt  with  cold  and  hot  KOH  (6).     (c)  By  the 
formation  of  the  silver  mirror  (§59,  1Gb).     (d)  By  its  action  as  an  alkali 
tartrate  in  preventing  precipitation  of  the  solutions  of  the  heavy  metals 
by  the  fixed  alkalis.     To  test  citric  acid  for  the  presence  of  tartaric  acid, 
add  about  one  cc.  of  ammonium  molybdate  solution  to  about  one  gram 
of  the  citric  acid;  then  two  or  three  drops  of  sulphuric  acid  and  warm 
on  the  water-bath.     The  presence  of  0.1  per  cent  or  more  of  tartaric  acid 
gives  a  blue  color  to  the  solution   (Crismer,  El.,  1891,   (3),  6,  23).     Add 
to  tartaric  acid  or  a  tartrate  a  little  ferrous  sulphate,  then  one  to  two  drops 
of  hydrogen  peroxide,  then  alkali — the  presence  of  tartaric  acid  will  be  in- 
dicated by  a  deep  violet  color  (Fenton,  Chem.  News,  43,  110).  Heat  tartaric 
acid  with  a  little  resorcin  and  concentrated  sulphuric  acid  in  a  porcelain 
dish  to  125-130°  C.     First,  red  streaks  will  appear,  then  the  whole  liquid 
turns  red.     Sensitiveness  0.01  mg.  tartaric  acid  (Mohler,  Bull.  Soc.  Chim. 
France   [3],  4,  1890,  728).     This  reaction  depends  upon  the  formation 
of  an  aldehyde  and  its  subsequent  condensation  with  resorcin. 

The  test  can  be  used  to  distinguish  between  citric  and  tartaric  acids. 
If  c*-naphthol  is  used  in  place  of  resorcin,  a  blue  liquid  turning  green  is 
the  result  (Pifieaua,  Chem.  Neivs,  9i,  179). 

9.  Estimation.— See  Philipps  (Z.,   1890,  29,  577);   Haas  (C.  C.,   1888,  1045); 
Heidenhain  (Z.,  1888,  27,  681). 


tARBON  MONOXIDE.  §226,  i. 

§226.  Carbon  monoxide.     CO  =  28.0  . 
C"0-",  C  =  0  . 

1.  Properties. — Carbon  monoxide,   carbonic  oxide,   formic  anhydride,   CO  ,   is  a 
colorless,  tasteless  gas.     Specific  gravity,  0.9678.     By  maintaining  a  pressure  of 
200  to  300  atmospheres  at  — 136°  and  then  reducing  the  pressure  to  50  atmos- 
pheres the  gas  becomes  a  colorless  transparent  liquid   (Wroblewski   and   Ols- 
zewski,  A.  €%.,  1884  (6),  1,  128).     It  is,  when  inhaled,  a  virulent  poison,  abstract- 
ing oxygen  from  the  blood  and  combining  with  the  haemoglobin.     It  burns  in 
the   air  with   a  pale   blue   flame   to   C02  ,   but   does   not   support   combustion. 
Mixed  with  air  in  suitable  proportions,  it  explodes  upon  ignition.     It  unites 
with  chlorine  in  the  sunlight  to  form  phosgene,  COCL  . 

2.  Occurrence. — In  combination  as  formic  acid  in  ants  and  in  nettles. 

3.  Formation. — (a)  By  the  incomplete  combustion  of  coal,  charcoal  or 
organic  material.     (b)  From  the  reduction  of  metallic  oxides  in  the  blast 
furnace  with  excess  of  charcoal:    Fe203  +  3C  =  2Fe  -j-  SCO.     (c)  By 
heating  sodium  sulphate  with  excess  of  charcoal  (LeBlanc's  soda  process) : 
Na2S04  +  4C  =  Na2S  +  4CO  .     See  also  Grimm  and  Eamdohr  (.4.,  1856, 
98,  127). 

4.  Preparation. — (a)  By  passing  steam  over  charcoal  at  a  white  heat 
(water  gas):   H20  +  C  =  CO  +  H2  (Naumann  and  Pistor,  B.,  1885,  18, 
164).     (b)  By    passing    C02    over    red    hot    charcoal,     (c)  By    heating 
K4Fe(CN)6  with  concentrated  H2S04:    K4Fe(CN)6  +  6H,S04  -f  6H20  - 
2K2S04  +  3(NH4)2S04  +  FeS04  +  6CO  .     With  dilute  acid  HCN  is  formed. 
(d)  By  heating  a  formate  with  concentrated  sulphuric  acid:    2KCH02  -f- 
H2S04  =  K2S04  +  2CO  +  2H20  .     (e)  By  heating  an  oxalate  with  con- 
centrated sulphuric  acid:    K2C204  +  2H0S04  =  K0S04  +  H,S04.H,0  -f 
CO  +  C02 . 

5.  Solubilities. — It  is  not  absorbed  by  KOH  or  Ca(OH)2   (distinction 
from  C02).     It  is  absorbed  by  charcoal,  cuprous  chloride,  and  by  several 
metals,  e.  g.,  K ,  Ag ,  and  An  . 

6.  Reactions. — It  is  an  energetic  reducing  agent.     Combines  with  moist 
fixed  alkalis  to  form  a  formate  (Froelich  and  Geuther,  A.,  1880,  202,  317). 
In  the  sunlight  it  combines  directly  with  chlorine  or  bromine.     It  is 
oxidized  to  C02  by  warming  with  K2Cr.>07  and  concentrated  H2S04  ;  also 
by  palladium  sponge  saturated  with  hydrogen,  and  in  presence  of  oxygen 
and  water  (Remsen  and  Keiser,  5.,  1884,  17,  83).    A  solution  of  PdCl2  is 
reduced  to  Pd  by  CO.     Reduces  iodine  pentoxide  I205  at  150°,  I205  -f~ 
5CO  =  I2  +  5C02. 

7.  Ignition. — When  heated  to  redness  with  Na  or  K,   carbon  and  an 
alkali   carbonate   are   formed.     Upon   ignition   of   metallic   oxides   in   an 
atmosphere  of  CO  a  reduction  of  the  metal  takes  place,  so  far  as  observed 
the  same  as  when  the  corresponding  metallic  forms  are  ignited  with  char- 
coal (Rodwell,  J.  C.,  1863,  16,  44). 


§227,  4,  a.  OXALIC  ACID.  263 

8.  Detection. — In  distinction  from  C02  by  its  failure  to  be  absorbed  by 
KOH  or  Ca(OH)2.     By  its  combustion  to  CO.,  and  detection  as  such.     By 
its  combination  with  hot  concentrated  KOH  to  form  a  formate.     By  its 
action  on  I205 .  It  is  detected  in  the  blood  by  the  absorption  spectrum  (Vogel, 
B.,  1878,  11,  235). 

9.  Estimation. — The  measured  volume  of  the  gas  is  brought   in  contact  with 
a  solution  of  cuprous    chloride    in    hydrochloric    acid     which    absorbs    the    CO 
(Thomas,  C.  N.,  1878,  37,  6).     The  gas  is  passed  through   a  U-tube   containing 
I  O5    heated   to    150°.     The   liberated   iodine   is    absorbed   in    RI   solution    and 
titrated  with  standard  Na2S2O3  solution.     Jour.  Am.  Chem.  Soc.,  1900,  22,  14. 

§227.  Oxalic  acid.     H2C204  =  90.016. 

0       0 

II       II  C02H 

H'2(C2)+60-"4,H  —  0  —  C  —  C  —  0  —  H     or     | 

C02H 

1.  Properties. — Absolute  oxalic  acid,  H2C204 ,  is  a  white,  amorphous  solid, 
which  may  be  sublimed  at  150°  with  only  partial  decomposition:  H2C2O4   = 
C02    +   CO    +   H,O  .     Crystallized   oxalic   acid,   H2C204,2H2O  ,    effloresces   very 
slowly  in  warm,  dry  air,  and  melts  in  its  water  of  crystallization  at  98°;  at 
which   temperature  the  liquid   soon   evaporates  to   the  absolute  acid.     Oxalic 
anhydride  is  not  formed. 

2.  Occurrence. — Found  in  many  plants  in  a  free  state  or  as  an  oxalate.     In 
sorrel  it  is  found  as  KHC2O4;  in  rhubarb  as  CaC204  .     As  ferrous  oxalate  in 
lignite  deposits;  as  ammonium  oxalate  in  guano. 

3.  Formation. — (a)  By   decomposition    of   cyanogen   with   water,   am- 
monium oxalate  being  one  of  the  products,      (fr)  By  the  oxidation  of 
glycol  with  nitric  acid,     (c)  By  heating  potassium  formate  above  400° 
(Merz  and  Weith,  B.,  1882,  15,  1507).     (d)  By  passing  C02  over  a  mixture 
of  sodium  and  sand  at  360°  (Drechsel,  BL,  1868,  10,  121). 

4.  Preparation. — (a)  By  action  of  nitric  acid  sp.  gr.  1.38  upon  sawdust, 
starch,  or  sugar.     By  the  continued  action  of  concentrated  nitric  acid, 
after  the  sugar  is  all  oxidized  to  oxalic  acid,  the  latter  is  farther  oxidized 
to   C02 .     (b)  By  heating  sawdust  with   KOH   or   NaOH .     Hydrogen  is 
evolved,  the  cellulose,  C6H1005 ,  heing  converted  into  oxalic  acid.     Under 
certain  conditions,  additional  products  are  formed.     It  is  also  formed  in 
the  oxidation  of  a  great  many  organic  compounds. 

CuH^On  +  12HN03  =  6HAO4  4.  12NO  +  11H2O 
3H2C204  +  2HN03  =  6C02  +  2NO  +  4H2O 
C6H1005  +  GKOH  +  H20  =  3K2C204  +  9H2 

Oxalates  are  formed:    a. — By  treating  the  oxide,  hydroxide,  or  car- 
bonate with  oxalic  acid.     In  this  manner  may  be  made  the  oxalates  of 


264  OXALIC  ACID.  §227,  4,  b. 

Pb,  Ag,  Hg',  Hg",  Sn".  Bi ,  Cu",  Cd ,  Zn ,  Al ,  Co ,  Ni ,  Mn ,  Fe",  Fe'", 
Cr'",  Ba ,  Sr ,  Ca ,  Mg  ,  Na  ,  K,  and  some  others. 

1). — By  adding  oxalic  acid  to  some  soluble  salt  of  the  metal.  In  this 
manner  the  above  oxalates  may  be  made,,  except  alkali,  magnesium, 
chromic,  ferric,  aluminum  and  stannic  oxalates,  which  are  not  precipitated. 
Antimonous  salts  are  precipitated,  but  the  precipitate  is  basic. 

c. — Alkali  oxalates  will  precipitate  the  same  solutions  as  oxalic  acid, 
but  many  of  the  precipitates  are  soluble  in  excess  of  the  alkali  oxalate, 
and,  as  a  rule,  the  salt  formed  is  a  double  one,  e.  g.,  AgNH4C,04  .  Ba ,  Ca 
and  Sr  are  well-defined  exceptions  to  this  rule — their  precipitates,  formed 
by  this  method,  being  normal  oxalates. 

d.—  Some  of  the  metals  when  finally  divided  are  attacked  by  oxalic  acid, 
hydrogen  being  evolved. 

5.  Solubilities.  —  Oxalic  acid  is  very  soluble  in  water  and  in  alcohol. 
Alkali  oxalates  are  freely  soluble  in  water,  as  is  also  chromic  oxalate. 
Nearly  all  other  metallic  oxalates  are  insoluble  in  water  or  only  sparingly 
soluble  (Luckow,  J.  C.,  1887,  52,  529). 

The  metallic  oxalates,  soluble  and  insoluble,  are  transposed  by  dilute 
sulphuric,  hydrochloric,  and  nitric  acids,  with  formation  of  oxalic  acid: 
CaC204  +  2HC1  =  CaCl2  +  H2C204 .  That  is :  the  precipitated  oxalates 
of  those  metals,  which  form  soluble  chlorides,  dissolve  in  dilute  hydro- 
chloric acid;  of  those  metals  which  form  soluble  sulphates,  in  dilute  sul- 
phuric acid;  and  all  precipitated  oxalates  dissolve  in  dilute  nitric  acid 

Acetic  acid  does  not  dissolve  precipitated  oxalates,  or  but  slightly. 
Certain  of  the  oxalates  dissolve,  to  some  extent,  in  oxalic  acid  (as  acid 
oxalates). 

6.  Reactions. — A. — With  metals  and  their  compounds. — Oxalic  acid  and 
soluble  oxalates  precipitate  solutions  of  many  of  the  metallic  salts.     With 
excess  of  the  alkali  oxalates  soluble  double  oxalates  of  the  heavy  metals 
are  frequently  formed  (4).     An  excess  of  alkali  oxalate  transposes  par- 
tially the  alkaline  earth  carbonates.     On  the  other  hand,  the  alkali  car- 
bonates in  excess  partially  transpose  the  alkaline  earth  oxalates  (Smith, 
J.  C.,  1877,  32,  245).     See  also  under  6b  of  the  respective  metals. 

Oxalic  acid  is  a  decided  reducing  agent,  being  converted  to  water  and 
carbonic  anhydride  (a),  and  the  metallic  oxalates  to  carbonates  and  carbonic 
anhydride  (&),  by  all  strong  oxidizing  agents. 

(a)     2H2C2O4  +  O2  =  2H2O  +  4C02 
(6)      2K2C204  +  02  =  2K2C03  +  2CO2 

-f • — Pb02  with  oxalic  acid  forms  lead  oxalate  and  C02 .  Oxalic  acid  has 
no  action  upon  Pb304 ,  but  reduces  it  quickly  in  presence  of  any  acici 
capable  of  changing  the  Pb304  to  PbO? , 


§227,  B  6.  OXALIC  ACID.  265 

2.  Oxalic  arid  or  ammonium  oxalate  boiled  in  the  sunlight  with  HgCl2 
gives  HgCl  and  CO,  \Gmd\ns  Hand-book,  9,  118]. 

3.— HaAs04  becomes  H  AsO, ,  and  CO,  is  evolved.  To  prove  that  Asv 
becomes  As"'.,  add  excess  of  potassium  hydroxide,  and  then  potassium  per- 
manganate. The  latter  will  be  quickly  decolored. 

4. — Bi,0-  becomes  bismuth  oxalate  and  C02  . 

5. — Mn"+n  becomes  Mn".  (That  is,  all  compounds  of  manganese  having 
more  than  two  bonds  are  reduced  to  the  dyad.)  In.  absence  of  other  free 
acid,  MnC,04  is  formed,  and  C02  is  given  off.  If  some  non-reducing  acid 
be  present,  such  as  H,S04 ,  it  unites  with  the  manganese,  and  all  of  the 
oxalic  acid  is  converted  into  C02 . 

6. — Co203  and  Co(OH)3  form  cobaltous  oxalate,  and  C02  is  evolved. 

7. — Ni203  and  Ni(OH)3  become  nickelous  oxalate,  and  CO,  is  evolved. 

6\— H2Cr04  is  reduced  to  chromic  oxalate,  and  C02  is  evolved. 

As  a  rule,  reducing  agents  have  no  action  on  oxalic  acid  at  ordinary 
temperatures.  By  fusion,  however,  a  few  metals,  K ,  Na ,  Mg ,  etc.,  reduce 
it  to  free  carbon. 

B. — With  non-metals  and  their  compounds. 

1.— HCN  ,  HCNS ,  H4Fe(CN)6 ,  and  H3Fe(CN)6  seem  to  be  without  action 
upon  oxalic  acid. 

2. — HN02  seems  to  have  no  action  upon  H2C204 .  With  HN03  ,  C02 , 
NO  ,  and  H,0  are  formed.  The  nitric  acid  should  be  concentrated.  Test 
for  the  C02  by  passing  the  gases  into  a  solution  of  BaCl2  containing  KOH  . 

3. — H3PO, ,  H3P03 ,  and  H3P04  do  not  act  upon  oxalic  acid. 

4. — Concentrated  sulphuric  acid,  with  a  gentle  heat,  decomposes  oxalic 
acidy  b}r  removing  the  elements  of  water  from  it,  with  effervescence  of 
carbon  dioxide  and  carbon  monoxide:  H2C204  -f-  H2S04  =  H2S04.H20  -f- 
C02  -f~  ^0  .  With  oxalates,  the  decomposition  generates  the  same  gases. 
Other  strong  dehydrating  agents  produce  the  same  result. 

The  effervescing  gases,  C02  and  CO  ,  give  the  reactions  for  carbonic  anhy- 
dride; also,  if  in  a  sufficient  quantity,  the  CO  will  burn  with  a  blue  flame, 
when  ignited. 

5. — With  chlorine,  hydrochloric  acid  is  formed  and  the  oxalic  acid 
becomes  C02  (Gmelin's  Hand-book,  9,  116).  This  reaction  takes  place 
more  readily  in  the  presence  of  KOH ,  forming  KC1  and  K2COS  .  HC10 
forms  C02  and  Cl .  If  the  oxalic  be  in  excess  HC1  is  formed.  The  action 
is  more  rapid  in  the  presence  of  a  fixed  alkali,  an  alkali  chloride  and 
carbonate  being  formed.  HC103  forms  C02  and  varying  proportions  of 
Cl  and  HC1 ,  a  high  degree  of  heat  and  excess  of  oxalic  acid  favoring 
the  production  of  HC1  (Calvert  and  Davieo,  J.  C.,  1850,  2,  193). 

6. — Bromine  decomposes  oxalic  acid   in.    alkaline   mixture,  forming  a 


266  OXALIC  ACID.  £227,  B  7. 

bromide  and  a  carbonate.  In  acid  mixture  a  similar  reaction  takes  place 
if  a  hot  saturated  solution  of  oxalic  acid  be  used  in  excess.  With  HBrO, , 
bromine  and  C02  are  formed;  with  excess  of  oxalic  acid  and  heat  hydro- 
bromic  acid  is  formed. 

7. — HI03  forms  C02  and  I .  With  mixtures  of  chlorates,  bromates,  and 
iodates,  the  chlorate  is  first  decomposed,,  then  the  bromate,  and  finally  the 
iodate  (Guyard,  J.  C.,  1879,  36,  593). 

7.  Ignition, — The  oxalates  are  all  dissociated  on  ignition.     Those  of 
the  metals  of  the  alkalis  and  alkaline  earths  are  resolved  at  an  incipient 
red  heat,  into  carbonates  and  carbon  monoxide  (a) — a  higher  temperature 
decomposing  the  alkaline  earth  carbonates.    The  oxalates  of  metals,  whose 
carbonates  are  easily  decomposed,  but  whose  oxides  are  stable,  are  re- 
solved into  oxides,  carbonic  anhydride,  and' carbon  monoxide  (6).     The 
oxalates  of  metals,  whose  oxides  are  decomposed  by  heat,  leave  the  metal, 
and  give  off  carbonic  anhydride  (c).     As  an  example  of  the  latter  class, 
silver  oxalate,  when  heated  before  the  blow-pipe,  decomposes  explosively, 
with  a  sudden  puffing  sound — a  test  for  oxalates : 

(a)     CaC204  =  CaC03  +  CO 

(6)     ZnC204  =  ZnO  +  C02  +  CO 

(c)      Agr2C204  =  2Ag  +  2C02 

8.  Detection. — (a)  By  warming  with  concentrated  sulphuric  acid  after 
decomposition  of  carbonates  with  dilute  sulphuric  acid;  showing  the  pres- 
ence of  C02  by  absorption  in  Ca(OH)2  or  in  a  solution  of  BaCL  alkaline 
with  KOH  ;  and  showing  the  presence  of  CO  by  its  combustibility.     (6)  In 
solution  by  precipitation  in  neutral,  alkaline,  or  acetic  acid  solution  by 
calcium  chloride,  and  solubility  of  the  precipitate  in  dilute  hydrochloric 
acid.     Frey  (Z.,  1894,  33,  533),  recommends  the  formation  of  a  zone  of 
precipitation.     To  the  HC1  solution  containing  BaCl2  and  CaCl2  he  adds 
carefully  a  solution  of  NaC2H302  and  watches  the  zone  of  contact,    (c)  Warm 
the  solution  with  dilute  H,S04  and  XMn04.     If  the  permanganate  is  not 
decolorized,  H2C204  is  absent;  if  decolorized  test  for  C02  with  Ca(OH)2. 

2KMnO4  +  5H2C2O4  +  3H2SO4  =  K2SO4  +  2MnSO4  +  8H2O  +  10CO2 . 

9.  Estimation.— (a)  It  is  precipitated  as  CaC2O4  ;    after  washing,  the  Ca  is 
determined  by   §188,  9,  from  which  the  oxalic    acid    is  calculated.     (6)   By  the 
amount  of  KMnO.  which  it  will  reduce,     (c)   By  measuring   the  amount  of 
evolved    when    it   is   oxidized   in    any     convenient    manner,    usually    by    MnO> . 
(d)   By  the  amount  of  gold  it  reduces  from  AuCl3 . 


§228,  4.  CARBON  DIOXIDE.  26? 

§228.  Carbon  dioxide.     C02  =  44.0  . 

(Carbonic  anhydride.) 
Carbonic  acid  (hypothetical).    H2C03  =  62.016  . 

0 

II 
CIV0~"2  and  H'2CIV0-"3 ,0  =  0  =  0  and  H  —  0  —  C  —  0  —  H. 

1.  Properties. — The  xjtccific  gravity  of  the  gas  CO2  is  1.52897  (Crafts,  C.  r.,  1888, 
106,  1G2);  of  the  liquid  at  —34°,  1.057   (Cailletet  and  Mathias,  C.  r.,  188G,  102, 
1202);  of  the  solid  (hammered),  slightly  under  1.2  (Landolt,  #.,  1884,  17,  309). 
Critical  temperature,  30.92°  (Andrews,  Trans.  Roy.  Soc.,  1869,  159,  583;  1876,  166, 
21).     It  is  a  heavy  colorless  gas;  which  at  low  temperatures,  +3°,  and  high 
pressure,  79  atmospheres,  may  be  condensed  to  a  clear  mobile  liquid;  and  upon 
further  cooling  this  becomes  a  snow-like  mass.     Liquid  C02  is  more  compres- 
sible than  other   liquids    (Natterer,  r/.,   1851,   59).     It  diffuses  through   porous 
plates  more  rapidly  than  oxygen  (Graham,  C.  N.,  1863,  8,  79).     Non-combustible 
and  a  non-supporter  of  combustion,  except  that  K ,  Na  and  Mg1  burn  in  the  gas 
forming  an  oxide  of  the  metal  and  free  carbon.     It  is  used  in  chemical  fire 
engines.     Non-poisonous  but  causes  suffocation  (drowning)  by  exclusion  of  air. 
It  is  taken  internally  without  injury  in  soda  water,  etc. 

Liquid  CO.,  is  insoluble  in  water  which  swims  on  the  surface.  It  mixes  with 
alcohol  and  ether.  It  dissolves  iodine  but  does  not  dissolve  phosphorus  or 
sulphur;  it  is  without  action  upon  K  or  Na .  A  spirit  thermometer  immersed 
in  the  liquid  registers  —75°  (Thilorier,  J.  /»*.,  1834,  3,  109).  Solid  CO,  at  767.3 
mm.  barometric  pressure  melts  at  — 77.94°  (Regnault,  A.  Ch.,  1849,  (3),  26,  257). 
When  the  solid  is  mixed  with  ether  it  gives  a  temperature  of  — 98.3°. 

2.  Occurrence. — In  a  free  state   in   the  air,   about  0.04  per  cent.     Found   in 
great  abundance  in  the  form  of  carbonates  in  the  earth's  crust;  e.  g.,  limestone, 
marble,  magnesite,  dolomite,  etc. 

3.  Formation. — (a)  By  burning  wood,  coal,   etc.,  in  the  air.     (b)  By 
burning  CO  .     (c)  By  the  reduction  of  many  metallic  oxides  upon  ignition 
with  charcoal,     (d)  During  fermentation  or  decay  of  organic  material. 
(e)  By  the  reaction  between  acids  and  carbonates. 

Liquid  C02  is  made  by  compressing  the  gas  with  pumps  at  a  reduced 
temperature. 

Solid  C02  is  made  by  allowing  the  liquid  to  escape  freely  into  woolen 
bags  and  then  compressing  in  wooden  moulds  (Landolt,  I.  c.). 

4.  Preparation. — CaC03  (chalk  or  marble)  in  small  lumps  is  treated  with 
hydrochloric  acid  in  a  Kipp's  gas  generating  apparatus.     The  gas  is  passed 
through  a  solution  of  NaHC03  to  remove  any  HC1  that  may  be  carried 
over,  and  then  dried  by  passing  through  a  tube  filled  with  fused  CaCl2 . 
It  is  also  prepared  on  a  large  scale  for  making  the  liquid  C02 ,  and  for 
use  in  sugar  factories  by  the  ignition  of  limestone :   CaC03  =  CaO  +  C02 . 

Preparation  of  Carbonates. — Na2C03  is  made  by  converting  NaCl  into 
Na-,S04 ,  by  treating  it  with  H2S04  ;  then  by  long  ignition  with  coal  and 
calcium  carbonate,  impure  sodium  carbonate  is  formed  (Leblanc's  process). 
Na.SO,  +  4C  +  CaCO;  =  CaS  +  4CO  +  Na,C03 


268  CARBOX  DIOXIDE.  §228,  5. 

It  is  separated  by  lixiviation  with  water,  and  farther  purified.  The 
other  method,  known  as  the  ammonia,,  or  Solvay's  process,  consists  in  pass- 
ing NH3  and  C02  into  a  concentrated  solution  of  NaCl  (a).  The  NaHC03 
is  converted  into  Na2C03  by  heat,  and  the  evolved  CO.,  used  over  again  (b). 
The  NH4C1  is  warmed  with  MgO  (c),  and  the  NH,  which  is  given  off  is 
used  over  again.  The  MgCl2  is  strongly  heated  (d)  and  the  MgO  is  used 
over  a'gain,  and  the  evolved  gas  sold  as  hydrochloric  acid.  This  continu- 
ous process  has  nearly  superseded  the  Leblanc  process. 

(a)     NaCl  +  NH3  +  H20  +  CO2  =  NaHC03  +  NH4C1 

(6)      2NaHCOs  +  heat  =  Na,C03  +  C02  +  H20 

(c)  2NH4C1  +  MgO  =  MgCL  +  2NH3  +  H20 

(d)  MgCL  +  H20  +  heat  ==  MgO  +  2HC1 

The  other  carbonates  are  mostly  made  from  the  sodium  salt  (6). 

5.  Solubilities. — C02    is    soluble    in    water,    forming    the    hypothetical 
H2C03 ,  which  reacts  acid  towards  litmus.     At  15°  one  volume  of  water 
absorbs  1.002  volumes  of  the  gas  (Bunsen,  A.,  1855,  93,  1).     It  is  rapidly 
absorbed  by  hydroxides  of  the  alkalis  and  of  the  alkaline  earths,  forming 
normal  or  acid  carbonates :    KOH  +  C02  ==  KHC03  or  2KOH  +  C02  - 
K2C03  +  H20  .     The  carbonates  of  the  alkalis  are  soluble  in  water  (acid 
alkali  carbonates   are   less   soluble   than   the   normal   carbonates),   other 
carbonates  are  insoluble  in  water  or  only  sparingly  soluble.     The  presence 
of  some  other  salts,  especially  ammonium  salts,  increases  the  solubility  oi 
carbonates,  notably  magnesium  carbonate  (§189,  5c).     Many  of  the  car- 
bonates are  soluble  in  water  saturated  with  C02  ;  forming  acid  carbonates 
of  variable  composition.     Boiling  removes  the  excess  of  C02 ,  causing  pre- 
cipitation of  the  carbonate. 

6.  Reactions. — Dry  carbon   dioxide  does   not  unite  with  dry  calcium 
oxide  at  ordinary  temperature  (Birnbaum  and  Maher,  B.,  1879,  12,  1547; 
Scheibler,  B.,  1886,  19,  1973).     Also  at  0°  no  reaction  takes  place  between 
dry  C02  and  dry  Na20 ,  but  at  400°  combination  takes  place  with  incan- 
descence (Beketoff,  El.,  1880,  (2),  34,  327). 

Carbonates  of  the  fixed  alkalis  precipitate  solutions  of  all  other  metallic 
salts:  with  antimony  the  precipitate  is  an  oxide;  with  tin,  aluminum, 
chromium,  and  ferricum  it  is  an  hydroxide;  with  silver,  mercurosum, 
cadmium,  ferrosum,  manganese,  barium,  strontium,  and  calcium  it  is  a  nor- 
mal carbonate;  with  other  metals  a  basic  carbonate,  except  that  mercuric 
chloride  forms  an  oxychlori.de.  Carbonic  acid  is  completely  displaced  by 
strong  acids,  for  example,  from  all  carbonates,  by  HC1  .  HC103 ,  HBr  ,  HBr03 , 
HI,  HI03,  H2C204,  HN03,  H3P04 ,  H,S04 ,  and  even  by  H2S ,  completely 
from  carbonates  of  the  first  four  groups,  incompletely  from  those  of  the 
fifth  and  sixth  groups  (Nandin  and  Montholon,  C.  r.,  1876,  33,  58). 

Ammonium  carbonate  precipitates  solutions  of  all  the  non-alkali  metals, 


§228,   6.  CARBON  DIOXIDE.  269 

chiefly  as  carbonates;' excepl  magnesium  salts  which  are  not  at  all  pre- 
cipitated, a  soluble  double  salt  being  at  once  formed  (separation  of  barium, 
strontium,  and  calcium  from  magnesium).  .With  salts  of  silver,  copper, 
cadmium,  cobalt,  nickel,  and  zinc  the  precipitate  is  redissolved  by  an 
excess  of  Hie  ammonium  carbonate. 

The  decomposition  of  carbonates  by  acids  is  usually  attended  by  marked 
effervescence  of  gaseous  C02  which  reddens  moist  litmus  paper:  Na2C03  + 
H2S04  r=  Na2S04  +  H20  +  C02 . 

With  normal  carbonates  in  cold  solution,  slight  additions  of  acid  (short 
of  a  saturation  of  half  the  base)  do  not  cause  effervescence,  because  acid 
carbonate  is  formed :  2Na2C03  +  H2S04  =  =  Na2S04  +  2NaHCOn  ;  and 
when  there  is  much  free  alkali  present  (as  in  testing  caustic  alkalis  for 
slight  admixtures  of  carbonate),  perhaps  no  effervescence  is  obtained. 
By  the  time  all  the  alkali  is  saturated  with  acid,  there  is  enough  water 
present  to  dissolve  the  little  quantity  of  .gas  set  free.  But  if  the  car- 
bonate solution  is  added  drop  by  drop  to  the  acid,  so  that  the  latter  is  con- 
stantly in  excess,  even  slight  traces  of  carbonate  give  notable  effervescence. 

The  effervescence  of  carbonic  acid  gas,  C02 ,  is  distinguished  from  that  of 
H2S  or  S02  by  the  gas  being  odorless,  from  that  of  N203  by  its  being  color- 
less and  odorless ;  from  all  others  by  the  effervescence  being  proportionally 
more  forcible.  It  should  be  remembered,  however,  that  C02  is  evolved 
(with  CO)  on  adding  strong  sulphuric  acid  to  oxalates  or  to  cyanates. 

On  passing  the  gas,  C02 ,  into  solution  of  calcium  hydroxide  (a);  or  of 
barium  hydroxide  (5);  or  into  solutions  of  calcium  or  barium  chloride, 
containing  much  ammonium  hydroxide  (c),  or  into  ammoniacal  solution 
of  lead  acetate  (d),  a  white  precipitate  or  turbidity  of  insoluble  carbonate 
is  obtained.  The  precipitate  may  be  obtained  by* decanting  the  gas  (one- 
half  heavier  than  air)  from  the  test-tube  in  which  it  is  liberated  into  a 
(wide)  test-tube,  containing  the  solution  to  be  precipitated;  but  the  opera- 
tion requires  a  little  perseverance,  with  repeated  generation  of  the  gas, 
owing  to  the  difficulty  of  displacing  the  air  by  pouring  into  so  narrow  a 
vessel.  The  result  is  controlled  better  by  generating  the  gas  in  a  large 
test-tube,  having  a  stopper  bearing  a  narrow  delivery- tube,  so  bent  as  to 
be  turned  down  into  the  solution  to  be  precipitated, 
(a)  C02  +  Ca(OH)2  =  CaC03  +  H2O 
(6)  CO,  +  BaCOH),  =  BaCO3  +  H2O 

(c)  CO2  +  CaCL  +  2NH4OH  =  CaCO,  +  2NH4C1  +  H2O 

(d)  C02  +  Pb20(C2H302)2  =  PbC03  +  Pb(C2H302)2 

The  solutions  of  calcium  and  barium  hydroxides  furnish  more  delicate 
tests  for  carbonic  anhydride  than  the  ammoniacal  solutions  of  calcium  and 
barium  chlorides,  but  less  delicate  than  lead  basic  acetate  solution.  The 
latter  is  so  rapidly  precipitated  by  atmospheric  carbonic  anhydride,  that 


270  fARBOy  DIOXIDE.  <22S.    T. 

it  cannot  be  preserved  in  bottles  partly  full  and  frequently  opened,  and 
cannot  be  diluted  clear,  unless  with  recently  boiled  water. 

Solutions  of  the  acid  carbonates  effervesce,  with  escape  of  C02 ,  on  boiling 
or  heating,  thus: 

iKHCO,  =  K2CO,  -f  H20  +  C02  .     (Gradually,  at  100°.) 

2NaHC03  =  Na;C03  -  H,0  -f  C02  .     (Gradually,  at  70°;  rapidly  at  90°  to  100°.) 
2NH4HC03  =  (NH4)2CO3  +  H50  -f  C02  .     (Begins  to  evolve  CO2  at  36°.) 
(NH4)4H2(C03),  =  2(NH4)2CO,  +  H,O  +  CO2  .     (Begins  at  49°.) 

7.  Ignition. — On  ignition,  the  normal  carbonates  of  the  metals  of  the 
fixed  alkalis  are  not  decomposed;  the  carbonates  of  barium  and  strontium 
are  dissociated  slowly,  at  white  heat,*  calcium  carbonate  at  a  full  red  heat, 
forming  the  oxide  and  C02 .     At  a  lower  temperature,  ignition  changes 
all  other  carbonates  to  the  oxide  and  C02 ,  except  that  the  carbonates  of 
silver  at  250°,  mercury,  and  some  of  the  rarer  metals  are  reduced  to  the 
metallic  state,  C02  and  oxygen  being  evolved.     Stannous  and  ferrous 
oxides  ignited  in  an  atmosphere  of  C02  are  changed  to  Sn02  and  Fe  0   . 
respectively,  with  evolution  of  CO  (Wagner,  Z.,  1879,  18,  559). 

8.  Detection. — Carbonates  are  detected:  (a)  By  the  sudden  effervescence 
when  treated  with  dilute  acids,     (b)  By.  the  precipitate  which  this  gas 
forms  with  solutions  of  Ca(OH)2,  Ba(OH), ,  or  Pb_0-C_H  0,)2.     If  but  a 
small  amount  of  carbonate  be  present,  the  mixture  must  be  warmed  to 
drive  the  C02  over  into  the  reagent  (6).     A  non-volatile  acid  as  H  SO 
H3P04  should  be  used,  as  a  volatile  acid  might  pass  over  with  the  C02  and 
prevent  the  formation  of  a  precipitate,     (c)  Phenolphthalein  detects  the 
normal  carbonate  in  solution  of  the  bicarbonate  (very  delicate).     Sodium 
bicarbonate  fails  to  give  a  precipitate  with  magnesium  sulphate  (distinc- 
tion from  Na2C03)  (Patein,  J.  Pharm.,  1892,  (5),  25,  448). 

To  detect  free  carbonic  acid  in  presence  of  bicarbonates,  a  solution  of 
1  part  of  rosolic  acid  in  500  parts  of  80  per  cent  alcohol  may  be  employed, 
to  which  barium  hydroxide  has  been  added  until  it  begins  to  acquire  a 
red  tinge.  If  0.5  cc.  of  this  rosolic  acid  solution  be  added  to  about  50  cc. 
of  the  water  to  be  tested — spring  water,  for  instance — the  liquid  will  be 
colorless,  or  at  most  faintly  yellowish  if  it  contains  free  carbonic  acid, 
whereas,  if  there  be  no  free  carbonic  acid,  but  only  double  salts,  it  will 
be  red  (Pettenkofer,  Dingl,  1875,  217,  158). 

Salzer  (Z.,  1881,  20,  227)  gives  a  test  for  free  carbonic  acid  or  bicar- 
bonates in  presence  of  carbonates,  founded  on  the  fact  that  tl 
ammonia  reaction  (§207,  6£)  does  not  take  place  in  presence  of  f  re- 
borne  acid  or  bicarbonates.     This  reaction  is  also  used  to  detect  the  presence 
of  fixed  alkali  hydroxides  in  the  fixed  alkali  carbonates.     In  presence  of  a 

*  Barium  carbonate  decomposes  at  1450°;  strontium  carbonate  &  1155°,  and  calcium 
Carbonate  at  825°, 


§£30,  1.  CYANOGEN— HYDROCYANIC  ACID. 

fixed  alkali  hydroxide  a  brown  precipitate  is  ol>i;iino<l  (Dobbin,  J.  ^oc.  Ind., 
1888,  7,  829). 

9.  Estimation. — (a)  By  decomposition  of  a  weighed  sample  with  a  known 
weight  of  anhydrous  borax  or  fused  microcosmic  salt  (NaPO?)  and  determining 
the  COo  by  loss  of  weight.  (6)  By  decomposition  of  the  weighed  sample  and 
collection  of  the  CO2  in  a  weighed  KOH  solution.  (c)  By  decomposition  with 
an  excess  of  a  standard  acid,  boiling  to  expel  the  CO,  and  titrating  the  excess  of 
acid,  (d)  Sodium  bicarbonate  may  be  estimated  by  titration  with  sodium 
hydroxide:  NaHCO3  +  NaOH  =  Na2CO3  +  H>O  .  The  first  excess  of  sodium 
hydroxide  beyond  the  reaction  gives  a  brown  precipitate  with  silver  nitrate 
(Lunge,  Z.  angew.,  1897,  169;  Bohlig,  Arch.  Pharm.,  1888,  226,  541). 

§229.  Cyanogen,     CN  =  26.01. 
N=C  — C  =N. 

A  colorless,  intensely  poisonous  gas;  specific  gravity,  1.8064  (Gay-Lussac,  Gilb., 
1816,  53,  145).  The  molecular  weight  shows  the  molecule  to  be  C2N2  .  At 
ordinary  atmospheric  pressure  it  liquifies  at  — 22°  (Drion,  J.,  1860,  41);  at  20° 
under  four  atmospheres  pressure  (Hofmann,  B.,  1870,  3,  658).  The  gas  has 
an  odor  of  bitter  almonds  and  burns  with  a  red  color  to  the  flame  forming 
C02  and  N  .  When  cooled  to  about  the  freezing  point  of  mercury  it  solidifies 
to  a  crystalline  ice-like  mass  (Hofmann,  I.  c.).  Critical  temperature,  124°  (De- 
war,  C.  2V.,  1885,  51,  27).  The  liquid  is  colorless,  mobile  and  a  non-conductor 
of  electricity.  It  occurs  in  the  gas  from  the  coke  ovens  (Bunsen  and  Playfair, 
J.  in:,  1847,  42,  145).  It  is  prepared:  (rt)  By  heating  the  cyanides  of  mercury, 
silver  or  gold:  Hg(CN)2  =  Hg  -f  C2N2  .  (ft)  By  the  dry  distillation  of  am- 
monium oxalate:  (NH4)2C,O4  =  4H2O  +  C2N2  .  (c)  By  fusing  KCN  with 
HgCL:  2KCN  +  Hg-Cl,  =  Hg  -f  2KC1  +  C2N2  .  (d)  By  heating  a  solution  of 
CuSO4  with  KCN  .  Half  of  the  CN  is  evolved  and  CuCN  is  formed.  If  the 
CuCN  be  heated  with  FeCl3  or  MnO,  and  HC2H3O2  ,  the  remainder  of  the 
CN  is  obtained.  The  gas  is  purified  by  absorption  with  aniline;  oxygen, 
nitrogen  and  carbon  dioxide  are  not  absorbed  (Jacquemin,  A.  Ch.,  1886,  (6),  6, 
140).  It  combines  with  Cl ,  Br  ,  I,  S,  P,  and  with  many  of  the  metals, 
reacting  very  much  like  the  halogens.  It  dissolves  in  water,  alcohol  and 
ether;  but  gradually  decomposes  with  formation  of  ammonium  oxalate  and 
carbonate  (Vauquelin,  A.  Ch.,  1823,  22,  132;  Buff  and  Hofmann,  A.,  1860,  "•  13, 
129).  At  500°  it  combines  with  hydrogen  to  form  HCN  (Berthelot,  BL,  1880, 
(2),  33,  2).  With  Zn  it  forms  Zn(CN)a  ,  rapidly  at  100°.  With  HC1  and  abso- 
lute alcohol  it  forms  oxalic  ether,  which  shows  cyanogen  to  be  the  nitrile  of 
oxalic  acid  (Pinner  and  Klein,  B.,  1878,  11,  1481).  With  solution  of  KOH, 
KCN  and  KCNO  are  formed:  C2N2  +  2KOH  =  KCN  +  KCNO  +  H2O  .  Com- 
pare the  reaction  with  chlorine  and  KOH  (§270). 


§230.  Hydrocyanic  acid.     HCN  =  27.018. 
H—  C  =  N. 

1.  Properties. — Hydrocyanic  acid  is  a  clear,  mobile  liquid,  boiling  at  26°.  At 
—15°  it  freezes  to  a  fibrous  crystalline  mass.  »S/>ecf/?c  gravity  at  19°,  0.697 
(Bleekrode,  I' me  Rnii.  Nor.,  1884,  37,  339).  It  burns  with  a  bluish-red  flame, 
forming  H2O  ,  CO2  and  N.  Its  index  of  refraction  is  much  less  than  that  of 
water  (Mascart,  C.  r.,  1878,  86,  321).  It  is  one  of  the  most  active  poisons 
known;  of  a  very  characteristic  odor,  somewhat  resembling  that  of  bitter 
almonds.  The  antidote  is  chlorine  or  ammonia  by  inhalation.  Its  water 
solution  decomposes  slowly,  forming  ammonium  formate:  scarcely  at  all  in 
tlit-  dark.  It  distils  readily  unchanged.  The  U.  S.  P.  solution  contains  two 
per  cent  of  HCN.  It  is  a  weak  acid,  scarcely  reddening  litmus;  its  salts  are 
partially  decomposed  by  CO*  .  The  free  acid  or  soluble  salts  when  warmed 


272  HYDROCYANIC  ACID.  §230,  2. 

with  dilute  alkalis  or  acids   (with  strong"  acids  in  the  cold)   becomes  formic 
acid  and  ammonia:  HCN  -f-  2H2O  =  HC02H  -f  NH3  . 

2.  Occurrence. — The  free  acid  dots  not  occur  in  nature,  but  in  combination 
in    the    kernels    of    bitter    almonds,    peaches,    apricots,    plums,    cherries    and 
quinces;  the  blossoms  of  the  peach,  sloe  and  mountain  ash;  the  lea\es  of  the 
peach,  cherry  laurel  and  Portugal  laurel;  the  3'oung  branches   of  the  peach; 
the  stem-bark  of  the  Portugal  laurel  and  mountain  ash;  and  the  roots  of  the 
last-named  tree,  when  soaked   in   water  for   a   time   and   then   distilled,  yield 
hydrocyanic  acid,  together  with  bitter-almond  oil.     Potassium  cyanide  appears 
in  the  deposits  of  blast  furnaces  for  the  smelting1  of  iron  ores. 

3.  Formation. — (a)    Decomposition   of  amygdaline  by   emulsine  and  distilla- 
tion.    (6)  By  the  action  of  the  electric  spark  on  a  mixture  of  acetylene  and 
nitrogen  (Berthelot,  </.,  1874,  113).     (c)  By  heating-  a  mixture  of  cyanogen  and 
hydrogen  (§229).     (d)  By  the  dry  distillation  of  ammonium  formate:  NH4CHO2 
=  HCN  -f-  2H2O  .     (e)  By  boiling  or  fusing  many  organic  compounds  contain- 
ing nitrogen  with  KOH  ,  forming  KCN   (Post  and  Huebner,  B.,  1872,  5,  408). 
(/)  By  decomposition  of  metallic  cyanides  with  mineral  acids,     (y)  By  heating 
chloroform  with   a  mixture   of    ammonium   and   potassium    hydroxides    (Hof- 
mann,  A.,  1867,  144,  116). 

4.  Preparation. —  (a)    Bj^  the   action   of  dilute   sulphuric  acid   on   potassium 
ferrocyanide:    2K4Fe(CN)8    +    3H2SO4    =    GHCN    -f    K2Fe2(CN)6    +    3X,S04  . 
(6)  By  action  of  acids  upon  metallic  cyanides.     '(€•)  By  the  action  of  sulphuric 
acid  upon  mercuric  cyanide  in  the  presence  of  metallic  iron:  Hg(CN).,  +  Fe  + 
H2S04  =  2HCN  +  FeS04  +  Hg  . 

Metallic  cyanides  are  prepared:  (a)  By  the  action  of  HCN  on  metallic 
hydroxides.  (6)  By  the  action  of  soluble  cyanides  on  metallic  salts,  (c)  By 
igniting  potassium  ferrocyanide:  KtFe(CN)6  =  4KCN  +  FeCa  +  N2  .  (d)  By 
heating  potassium  ferrocj^anide  with  potassium  carbonate.  If  prepared  in 
this  manner  it  contains  some  cyanate:  K4Fe(CN)6  +  K.,CO3  =  5KCN  +  KCNO 
-f  Fe  +  C02  . 

5.  Solubilities. — Hydroc3ranic  acid  is  soluble  in  water,  alcohol  and  ether  in 
all  proportions.     A  mixture  of  equal  parts  acid  and  water  increases  in  tem- 
perature from  14°  to  22.5°;  it  also  increases   slightly   in  volume    (Bussy    and 
Buignet,  A.  Ch.,  1865,  (4),  4,  4). 

The  cyanides  of  the  alkali  metals,  alkaline  earth  metals,  and  mercuric 
cyanide,  are  soluble  in  water,  barium  cyanide  being  but  sparingly  soluble. 
The  solutions  are  alkaline  to  test-paper.  The  other  metallic  cyanides  are 
insoluble  in  water.  Many  of  these  dissolve  in  solutions  of  alkali  cyanides, 
by  combination,  as  double  metallic  cyanide*. 

Pb  ,  Hg-,  As,  Sb  ,  Sn  ,  Bi  and  Cd  are  dissolved  by  KCN  with  absorption 
of  oxygen.  Cu  ,  Al ,  Fe  (by  H  or  CO),  Co,  Ni  ,  Zn^and  Mg  with  evolution 
of  hydrogen:  2Cu  +  2KCN  *+  2H,0  =  2CuCN  +  2KOH  -f-  H2  .  Iron  or  steel 
wire  are  not  attacked  (Goyder,  C.  N.,  1894,  69,  202,  268  and  280). 

6.  Reactions. — There  are  two  classes  of  double  cyanides,  both  of  which  are 
formed   when  a  cyanide  is  precipitated   bv  an   alkali   cyanide,   and   redissolvrd 
by  excess  of  the  precipitant:  HgCL  +  2KCN  =  Hg(CN),   +  2KC1;  and  with 
excess  of  KCN:  Hg>(CN)2  -f  2KCN  =  (KCN)2Hg(CN)2  . 

Class  I.  Double  cyanides  which  are  not  affected  by  alkali  hydroxides,  but  are  decom- 
posed when  treated  with  dilute  acids:  '  KCN)2Hg(CN)2  +  2HC1  =  Hg(CN)2  +  2KC1 
-4-  LHCN  .  These  closely  resemble  the  double  iodides  (potassium  mercuric),  and 
the  double  sulphides  or  thiosalts  (§69,  5c  and  6e).  The  most  frequently  occurring 
of  the  double  cyanides  of  this  class,  which  dissolve  in  water,  are  given  below: 

Potassium  (or  sodium)  zinc  cyanide,  K2Zn(CN)4  or  (KCN)2Zn(CN)2 . 

Potassium  (or  sodium)  nickel  cyanide,  K2Ni(CN)4  or  (KCN)2Ni(CN)2 . 

Potassium  (or  sodium)  copper  cyanide,  K2Cu(CN)3  or  (KCN)2CuCN . 

Potassium  cadmium  cyanide,  K2Cd(CN)4  or  (KCN)oCd(CN)2 . 

Potassium  (sodium  or  ammonium)  silver  cyanide,  KAg(CN)2  or  KCNAgCN  . 

Potassium  (or  sodium)  mercuric  cyanide,  K2Hg(CN)4  or  (KCN)2Hg(CN)2  • 

Potassium  (or  sodium)  auric  cyanide,  KAu(CN)4  or  KCNAu(CN)3 . 

Class  II.  Double  cyanides  ivhich,  as  precipitates,  are  transposed  by  alkali  hydroxides, 
in  dUute  solution  ('/),  and  are  transposed,  without  decomposition,  by  dilute  acids  (6). 
In  these  double  cyanides,  as  potassium  ferrous  cyanide,  K4Fe(CN)6,  the  whole  of 


§230,  6.  HYDROCYANIC  ACID.  273 

the  cyanogen  appears  to  form  a  new  compound  radical  with  that  metal  whose  single 
cyanide  is  insoluble  in  water;  thus,  Fe(CN)e  as  "ferrocyanogen,"  giving  K4Fe(CN)e 
as  "potassium  ferrocyanide  "  (for  the  potassium  ferrous  cyanide).  These  more 
stable  double  cyanides  or  "ferrocyanides,"  etc.,  correspond  to  the  platinic  double 
chlorides  or  "  chloroplatinates  "  (§74,  5-),  and  the  palladium  double  chlorides,  or 
chloropalladiates  (§106,  5';).  The  most  frequently  occurring  of  the  double  cyanides 
of  this  class,  which  are  soluble  in  water,  are  given  below. 

(a)     Cu2Fe(CN)6  +  4KOH  =  2Cu(OH)2  +  K4Fe(CN)6 
(&)     K4Fe(CN)6  +  2H2S04  =  2K,SO4  +  H4Fe(CN)6 
2K3Fe(CN)6  +  3H2SO4  =  3K2SO4  +  2H3Fe(CN)6 

Alkali  ferrocyanides,  as  K4Fe"(CN)6  ,  potassium  ferrocyanide. 

Ferricyanides,  as  K3Fe'"(CN)6  ,  potassium  ferricyanide. 

Cobalticyanid.es,  as  K3Co'"(CN)<; ,  potassium  cobal  icyani;le. 

Man  gam  cyanides,  as  K3Mn'"(CN)6 ,  potas-ium  manganicyanide. 

Chromicyanides,  as  K3fCr"')(CN)6  ,  potassium  chromicyanide. 

The  easily  decomposed  double  cyanides  of  Class  I  are,  like  the  single  cyan- 
ides, intensely  poisonous.  The  difficultly  decomposed  double  cyanides  of 
Class  II.  are  not  poisonous. 

The  Single  Cyanides  are  transposed  by  the  stronger  mineral  acids,  more 
or  loss  readily,  with  liberation  of  hydrocyanic  acid,  HCN,  effervescing  from 
concentrated  or  hot  solutions,  remaining  dissolved  in  cold  and  dilute  solu- 
tions. Mercuric  cyanide  furnishes  HCN  by  action  of  H2S ,  not  by  other 
acids.  The  cyanides  of  the  alkali  and  alkaline  earth  metals  are  decomposed 
by  all  acids— even  the  carbonic  acid  of  the  air— and  exhale  the  odor  of 
hydrocyanic  acid.  Solution  of  silver  nitrate  precipitates,  from  solutions 
of  cyanides  or  of  hydrocyanic  acid  (not  from  mercuric  cyanide)  silver 
cyanide,  AgCN ,  white,  insoluble  in  dilute  nitric  acid,  soluble  in  ammonium 
hydroxide,  in  hot  ammonium  carbonate,  in  potassium  cyanide,  and  in 
thiosulphates — uniform  with  silver  chloride.  Cold  strong  hydrochloric 
acid  decomposes  it  with  evolution  and  odor  of  hydrocyanic  acid  (recogni- 
tion from  chloride);  and  when  well  washed,  and  then  gently  ignited,  it  does 
not  melt,  but  leaves  metallic  silver,  soluble  in  dilute  nitric  acid,  and  pre- 
cipitable  as  chloride  (distinction  and  means  of  separation  from  chloride). 

Solution  of  mercurous  nitrate,  with  cyanides  or  hydrocyanic  acid,  is 
resolved  into  metallic  mercury,  as  a  gray  precipitate,  and  mercuric  cyanide 
and  nitrate,  in  solution.  Salts  of  copper  react,  as  stated  in  §77,  66;  salts 
of  lead,  as  stated  in  §57,  6&. 

Ferrous  salts,  added  to  saturation,  precipitate  from  solutions  of  cyan- 
ides, not  from  hydrocyanic  acid,  ferrous  cyanide,  Fe(CN)2 ,  white,  if  free 
from  the  ferric  hydroxide  formed  "by  admixture  of  ferric  salt,  and,  with 
the  same  condition,  soluble  in  excess  of  the  cyanide,  as  (with  potassium 
cyanide),  (KCN)4Fe(CN)2  ==  K4Fe(CN)6 ,  potassium  ferrocyanide  (a).  On 
acidulating  this  solution,  it  gives  the  blue  precipitates  with  ferric  salts  (b)  : 
(a)  2KCN  +  FeSO4  =  Fe(CN)2  +  K2SO4 

Pe(CN)2  +  4KCN  =  K4Fe(CN)6 
(6)     3K4Fe(CN)6  +  4FeCl3  =  Fe4(Fe(CN)fl)»  +  12KC1 

This  production  of  the  blue  ferric  ferrocyanide  is  made  a  delicate  test  for 


274  HYDROCYANIC  ACID.  §230,  V. 

hydrocyanic  acid,  as  follows:  A  little  potassium  hydroxide  and  ferrous 
sulphate  are  added,  the  mixture  digested  warm  for  a  short  time;  then 
a  very  little  ferric  chloride  is  added,  and  the  whole  slightly  acidulated 
(so  as  to  dissolve  all  the  ferrous  and  ferric  hydroxides),  when  Prus- 
sian blue  will  appear,  if  hydrocyanic  acid  was  present  (Link  and 
Moeckel,  Z.,  1878,  17,  456.  For  identification  of  traces  of  hydrocyanic 
acid  (less  than  0.00002  g.  in  1  c.c.)  add  two  drops  of  10%  solution  of 
sodium  hydroxide,  evaporate  almost  to  dryness,  cool,  add  one  drop  of  a 
2%  solution  of  ferric  sulphate  and  allow  to  stand  in  the  cold  for  10-15 
minutes.  Heat  gently  with  two  or  three  drops  of  strong  hydrochloric 
acid  and  cool.  The  undiluted  blue-green  solution  shows,  when  carefully 
diluted,  a  blue  color.  (G.  Druce  Lander  und  Walden,  nach  Leitschr.  J. 
Unters.  d.  Nahrungs  u.  Genussm.  23  [1912]  399.) 

Solution  of  nitrophenic  acid,  picric  acid,  C6H2(N02)3OH ,  added,  in  a 
small  quantity,  to  a  neutralized  solution  of  cyanides  of  alkali  meta:s,  on 
boiling(  and  standing),  gives  a  blood-red  color,  due  to  picrocyanate  (as 
KC8H4N506).  This  test  is  very  delicate,  but  not  very  distinctive,  a^-  var- 
ious reducing  agents  give  red  products  with  nitrophenic  acid  (\ogel, 
C.  N.,  1884,  50,  270). 

The  fixed  alkali  hydroxides,  in  boiling  solution,  strongly  alkaline,  g  mdu- 
ally  decompose  the  cyanides  with  production  of  ammonia  and  formate: 
HCN  +  KOH  +  H20  =  KCH02  +NH, .  Ferrocyanides  and  ferricya aides 
finally  yield  the  same  products.  Dilute  alkalis,  not  heated,  transpose,  as 
by  equation  a,  class  II  above. 

Cyanides  are  strong  reducing  agents.  The  action  is  not  so  mark  3d  in 
solution  as  in  state  of  fusion  (7).  Permanganates  are  reduced  by  cyan- 
ides, and  cupric  hydroxide  in  alkaline  solution  forms  Cu'.  Solvtions 
of  cyanides  on  exposure  to  the  air  take  up  some  oxygen  with  formation  of 
a  cyanate:  2KCN  -(-  02  =  2KCNO  .  Commercial  potassium  cyanide  a  ways 
contains  some  potassium  cyanate.  By  warm  digestion  of  a  cyanide  with 
sulphur  or  with  yellow  ammonium  sulphide  a  thiocyanatc  is  formed  (8). 
Hydrocyanic  acid  reduces  Pb02 ,  forming  Pb(CN)2  and  CN  :  PbO,  +  ;;HCN 
=  Pb(CN),  +  C2N2  +  2H20  (Liebig,  A.,  1838,  25,  3).  With  HCI:  and 
H202  oxamide  is  formed  (Altfield,  J.  C.,  1863,  16,  94).  Chlorine  forms 
with  hydrocyanic  acid  a  cyanogen  chloride  (Serullas,  A.  Ch.,  182cc,  38, 
370);  with  iodine  the  reaction  is  not  so  marked,  but  a  similar  product  is 
formed  (Meyer,  P.,  1887,  20,  III,  704).  Concentrated  sulphuric  acid 
decomposes  all  cyanides. 

7.  Ignition, — By  fusion  with  fixed  alkalis,  cyanides  and  all  compounds 
containing  cyanogen  yield  ammonia.  In  state  of  fusion  cyanides  are  very 
efficient  reagents  for  reduction  of  metals  from  their  oxides  or  sulphides. 


§231.  HYDROFERROCYANIC   ACID.  275 

to  il  e  metallic  state  (§69,  !).      The  cyiiiuttcs  or  thiocyanates  formed  in  the 
reaction  are  not  readily  decomposed  by  heat  alone. 

8.  Detection. — Cyanides  may  he  detected:    (a)   By  the  odor  of  the  free 
arid    upon   decomposition  of   I  he  cyanide  with   acids.     This  test  must  be 
appl  }d  with  extreme  caution  as  the  evolved  HCN  or  CN  is  very  poisonous. 
(/>)   I  v  formation  of  a  ferrocyanide  and  its  reaction  with   ferric  salts,  as 
dese-  :hed  in  G.     (c)  The  production  of  the  red  ferric  thiocyanate  is  a  test 
for  i  ifdrocyanic  acid,  more  delicate  than  formation  of  ferrocyanide.     By 
warn,  digestion  this  reaction  occurs:   2KCN  -f-  S2  =  2KCNS  ;  or: 

2(NH4)2S4  +  4HCN  =  4NH.CNS  +  2H.S  +  S2 

To  the  material  in  an  evaporating-dish,  add  one  or  two  drops  of  yellow 
ammonium  sulphide,  and  digest  on  the  water-hath  until  the  mixture  is 
colorless,  and  free  from  sulphide.  Slightly  acidulate  with  hydrochloric 
acid  (which  should  not  liberate  H2S),  and  add  a  drop  of  'solution  of  ferric 
chlonde;  the  blood-red  solution  of  ferric  thiocyanate  will  appear,  if  hydro- 
cyan  c  acid  was  present  (Link  and  Moeckel,  /.  c.). 

(d:  Link  and  Moeckel  also  recommend  the  following  test  for  cyanides, 
delic  ite  to  1-3,000,000.  Saturate  a  fdter  paper  with  a  four  per  cent 
alcololic  solution  of  guaiac;  allow  the  alcohol  to  evaporate;  then  moisten 
the  "paper  with  a  one-fourth  per  cent  solution  of  copper  sulphate,  and 
allow  the  unknown  solution  to  trickle  over  this  test  paper.  A  deep  blue 
cokn  indicates  the  presence  of  a  cyanide. 

TV  detect  cyanides  in  presence  of  ferri-  and  ferrocyanides  it  is  directed 
to  add  tartaric  acid  and,  in  a  distilling  flask,  pass  a  current  of  carbon 
dioxile,  warming  not  above  (>0°.  Test  the  distillate  by  the  methods 
given  above.  Ferro-  and  ferricyanides  do  not  yield  HCN  under  80°  (Hilger 
and  Tamba,  Z.,  1891,  30,  529;  also  Taylor,  C.  N.,  1884,  50,  227). 

9.  Estimation. — (a)  The  nearly  neutral  solution  of  cyanide  is  titrated  with 
standard   silver   nitrate.     No   precipitate   occurs   as   long   as   two   molecules   of 
alkali   cyanide   are   present   to   one   of   silver   nitrate.     Soluble   AgCN,KCN   is 
form  id.     As   soon   as  the   alkali   cyanide   is  all   used   in  the   formation   of   the 
doub  e  cyanide,  the  next  molecule  of  silver  nitrate  decomposes  a  molecule  of 
tJie   double  salt,  forming-  two  molecules  of  insoluble  silver  cyanide;   giving  a 
white  precipitate  for  the  end  reaction.     Chlorides  do  not  interfere  (Liebig,  A., 
1851,  77,  102).     (6)  By  titratioii  with  a  standard  solution  of  HgCl2  ,  applicable 
in  presence  of  cyanates  and  thiocyanates  (Hannay,  J.  C.,  1878,  33,  245). 

§231.  Hydroferrocyanic  acid.     H4Fe(CN)0  =  215.932. 
H'4Fe"(CN)-'0 . 

Absolute  hydroferrocyanic  acid  (§230,  0,  Class  IT.),  is  a  white  solid,  freely 
soluble  in  water  and  in  alcohol.  The  solution  is  strongly  acid  to  test-paper, 
and  decomposes  carbonates,  with  effervescence,  and  acetates.  It  is  non-volatile, 
but  absorbs  oxygen  from  the  air,  more  rapidly  when  heated,  evolving  hydro- 
cyanic acid  and"  depositing  Prussian  blue:  7H4Fe(CN)0  +  O.,  =  Fe4  (Fe(CN)6)3 
-f  2H2O  +  24HCN  . 

Potassium  ferrocyanide  is  the  usual  starting-  point  in  the  preparation  of  the 


276 


HYDROFERROCYANIC  ACID. 


§231. 


free  acid  or  any  of  the  salts.  It  is  prepared  by  fusing  together  in  an  iron 
kettle  nitrogenous  animal  matter  (blood,  hair,  horn,  hoof,  etc.),  commercial 
potash  (KOH),  and  scrap  iron.  The  ferrocyanide  is  formed  when  this  mass  is 
digested  with  water.  The  nitrate  is  evaporated  to  crystallization  (lemon-yellow 
prism),  soluble  in  four  parts  of  water. 

Hydroferrocyanic  acid  is  formed  by  transposition  of  metallic  ferrocyanides 
in  solution,  with  strong1  acids  (a).  When  the  solution  is  heated,  hydrocyanic 
acid  is  evolved;  in  the  case  of  an  alkali  ferrocyanide,  without  absorption  of 
oxygen  (&).  Potassium  ferrocyanide  and  sulphuric  acid  are  usually  employed 
for  preparation  of  hydrocyanic  acid  (c) : 

(a)     K4Fe(CN)6  +  2H2S04  =  2K2S04  +  H4Fe(CN)6 

(6)     3H4Fe(CN)6  +  K4Fe(CN)6  =  2K2FeFe(CN)G  +  12HCN 

(c)     2K4Fe(CN)G  +  3H2S04  =  3K2S04  +  K2FeFe(CN)6  +  6HCN 

The  ferrocyanides  of  the  alkali  metals,  strontium,  calcium  and  magnesium, 
are  freely  soluble  in  water;  of  barium,  sparing^  soluble;  of  the  other  metals, 
insoluble  in  water.  There  are  double  ferrocya Hides;  soluble  and  insoluble;  that 
of  barium  and  potassium  is  soluble,  but  potassium  calcium  ferrocya&ide  is  in- 
soluble. The  most  of  the  ferrocyanides  of  a  heavy  metal  and  an  alkali  metal 
are  insoluble.  Potassium  and  sodium  ferrocyanides  are  precipitated  from  their 
water  solutions  by  alcohol  (distinction  from  ferricyanides). 

The  soluble  ferrocyanides  are  yellowish  in  solution  and  in  crystals,  white 
when  anhydrous.  The  insoluble  ferrocyanides  have  marked  and  very  diverse 
colors,  as  seen  below. 

Solutions  of  alkali  ferrocyanides,  as  K4Fe(CN)6  ,  give,  with  soluble  salts  of: 

Aluminum,    a    white    precipitate,  A1(OH)3  and  Fe(CN)2   (formed  slowly). 


Antimony  a  white 
Bismuth,  a  white 
Cadmium,  a  white 
Calcium,  a  white 
Chromium,  no 
Cobalt,  a  green,  then  gray 
Copper,  a  red-brown 
Gold,  no 

Iron  (Fe"),  w^hite,  then  blue 
Iron  (Fe'"),  a  deep  blue 
Lead,  a  white 
Magnesium,  a  white 

a  yellow-white 

Manganese,  a  white 
Mercury  (Kg-'),  a  white 
Mercury  (Hg"),  a  white 


Sb4[Fe(CN)0]3.25H20. 
Bi4(Fe(CN)6)3. 
Cd2Fe(CN)G  (soluble  in  HC1). 
K2CaFe(CN)0  . 

Co2Fe(CN)6  . 
Cu2Fe(CN)6  . 

K2FeFe(CN)6  . 

Fe4(Fe(CN)6)3. 

Pb2Fe(CN)6  . 

(NH4)2MgFe(CN)6  (in  presence  of  NH4OH) 

K2MgFe(CN)0  (only  in  concentrated  solu- 
tion). 

Mn2Fe(CN)6  (soluble  in  HC1). 

Hg4Fe(CN)6  (gelatinous). 

Hg2Fe(CN)0  ,  turning  to  Hg(CN)2  and 
Fe3(Fe(CN)6)2,  blue. 


Molybdenum,  a  brown 
Nickel,  a  greenish-white 
Silver,  a  white 
Tin  (Sn"  and  Sniv),  white 
Uranium  (uranous),  brown 
TJranium  (uranyl),  red-brown 
Zinc,  a  white,  gelatinous 


(slowly  turning  blue). 


Ni2Fe(CN)6 

Ag4Fe(ClSr)8 

(gelatinous). 

UFe(CN)6  . 

(U02)2Fe(CN)6. 

Zn2Fe(CN)G  . 

See  Wyrouboff  (A.  Ch.,  1876  (5),  8,  444;  and  1877,  (5),  10,  409). 
Insoluble  ferrocyanides  are  transposed  by  alkalis  (§230,  6,  Class  II.) 
It  will  be  observed  (§230,  6)  that  ferrocyanides  are  ferrous  combinations,  while 
ferricyanides  are  ferric  combinations.     And,  although  ferrocyanides  are  far  less 
easily  oxidized  than  simple  ferrous  salts,  being  stable  in  the  air,  they  are 


§232.  HYDROFERRICYANIC  ACID.  277 

nevertheless    reducing-    agents,    of    moderate    power:    2K4Fe(CN)8     +    CL    = 

2K3Fe(CN)(i  +  2KC1  . 

PbO.  with  sulphuric  acid  forms  Pb"  and  H3Fe(CN)0  . 

Ag'  \vitli  fixed- alkali  forms  an  alkali  ferricyanidc  and  metallic  silver. 

CrVi  with  phosphoric  acid,  gives  Cr'"  and  H3Fe(CN)0    (Schonbein,  J.  pr.,  1840, 

20,  145). 

Co'"  with  phosphoric  acid  forms  Co"  and  H8Fe(CN)0  . 
Ni'"  with  acetic  acid  gives  Ni"  and  H3Fe(CN)«  . 
Mn02  with  phosphoric  acid  gives  Mn"  and  H3Fe(CN)0  . 
Mnvn    forms    with    potassium    hydroxide    Mn(X    and    potassium    ferricyanide. 

\Vith  sulphuric  acid,  manganous  sulphate  and  hydroferricyanic  acid. 
Ferricyanides  when   boiled   with   NH4OH   give  ferrocyanides    (Playfair,  J.    C., 

1857,  9,  128). 
HNO.,  forms  first  hydroferricyanic  acid,  then  hydronitroferricyanic  acid  and 

NO. 
HNO3   forms  hydroferricyanic  acid,  and  then  hydronitroferricyanic  acid,   NO 

being  evolved. 
Cl  forms  first  hydroferricyanic  and  hydrochloric  acids.     Excess  of  chlorine  to 

be  avoided  in  preparation  of  ferricyanides. 
HC103  forms  hydroferricyanic  and  hydrochloric  acids. 
Br  forms  hydroferricj^anic  and  hj^drobromic  acids. 
HBrO3  forms  hydroferricyanic  and  hydrobromic  acids. 

I ,  iodine  is  decolored  by  potassium  ferrocyanide,  and  some  potassium  ferri- 
cyanide and  potassium  iodide  are  formed.     The  action  is  slow  and  never 

complete  (Cmdin's  Hand-book,  7,  459). 
HI03  forms  hydroferricyanic  acid  and  free  iodine. 

In  analysis,  soluble  ferrocyanides  are  recognized  by  their  reactions  with 
ferrous  and  ferric  salts  and  ccpper  salts  (see  6ft,  §126  and  §77).  Separated 
from  ferricyanide,  by  insolubility  of  alkali  salt  in  alcohol.  Separation  of  hydro- 
ferrocyanic  acid  from  hydroferricyanic  acid  according  to  Ph.  E.  Browning  and 
H.  E.  Palmer  (Zeitschr.  f.  anorg.  chem.,  54,  315,  nach.  Zeitschr.  f.  anal,  chem., 
60  (1911)  771).  Acidify  5-10  c.c.  of  the  solution  which  is  to  be  tested  with  acetic 
or  hydrochloric  acid  and  add  a  solution  of  a  thorium  salt.  Thorium  ferrocyanide 
will  be  precipitated.  Shake  up  with  finely  divided  asbestos,  filter,  wash  the  pre- 
cipitate, add  sodium  hydroxide  to  same  and  in  this  filtrate  test  for  ferrocyanide. 

To  the  filtrate  from  the  thorium  ferrocyanide  add  a  solution  of  a  cadmium  salt. 

Cadmium  ferricyanide  will  be  precipitated  and  is  treated  like  thethorium  precipitate. 

Ferrocyanides  are  estimated  in  solution  with  sulphuric  acid  by  titrating  with 

standard  KMnO4 .     Also  by  precipitation  with    CuSOt    either  for  gravimetric 

determination  or  volumetrically,  using  a  ferric  salt  as  an  external  indicator. 

§232.  Hydroferricyanic  acid.     H3Fe(CN)6  =  214.924. 
H'3Fe"'(CN)-'6. 

Absolute  hydroferricyanic  acid,  H3Fe(CN)6  ,  is  a  non- volatile,  crystallizable 
solid,  readily  soluble  in  water,  with  a  brownish  color,  and  an  acid  reaction  to 
test-paper.  It  is  decomposed  by  a  slight  elevation  of  temperature.  In  the 
transposition  of  most  ferricyanides,  by  sulphuric  or  other  acid,  the  hydro- 
ferricyanic acid  radical  is  broken  up. 

Potassium  ferricyanide  is  the  usual  starting  point  in  the  preparation  of  most 
ferricyanides.  It  is  prepared  by  passing  chlorine  into  a  cold  solution  of 
K4FeiCN)f;  until  a  few  drops  of  the  liquid  gives  a  brownish  color,  but  no  pre- 
cipitate with  a  ferric  salt.  The  solution  is  evaporated  to  crystallization  and  the 
salt  repeatedly  recrystallized  from  water  as  large  red  prismatic  crystals, 
very  soluble  in  water,  freely  soluble  in  alcohol  (distinction  from  KjFe(CN)6). 
The  free  acid  is  made  by  adding  to  a  cold  saturated  solution  of  K3Fe(CN)6 
three  volumes  of  concentrated  HC1  and  drying  the  precipitate  which  forms, 
in  a  vacuum  (Joannis,  C.  r.,  1882,  94,  449,  541  and  725)  lustrous,  brown  sh- 
green  needles,  very  soluble  in  water  and  alcohol,  insoluble  in  ether. 

The  ferricyanides  of  the  metals  of  the  alkalis  and  alkaline  earths  are  soluble 
in  water;  those  of  most  of  the  other  metals  are  insoluble  or  sparingly  soluble. 
The  soluble  ferricyanides  have  a  red  color,  both  in  crystals  and  solution;  those 
insoluble  have  different,  strongly  marked  colors.  Potassium  and  sodium  ferri- 


278  H  Y  DROP  ERHI  CYAN  1C  ACID.  §232. 

cyanides  are  but  slightly,  or  not  at  all,  precipitated  from  their  water  solutions 
by  alcohol  (separation  from  ferrocyanides). 

Ferricyanides  are  not  easily  decomposed  by  dilute  acids;    but  alkali  hydrox- 
ides, either  transpose  them  or  decompose  their  radicals  (§230,  6). 

Solutions  of  metallic  ferricyanides  give,  with  soluble  salts  of: 
Aluminum,  no  precipitate. 
Antimony,  no  precipitate. 

Bismuth,  light-brown  precipitate,  BiFe(CN)0  ,  insoluble  in  HC1  . 
Cadmium,  yellow  precipitate,  Cd,[-Fe(CN)<l]s  ,  soluble  in  acids  and  in  ammo- 

nium hydroxide. 
Chromium,  no  precipitate. 
Cobalt,  brown-red  precipitate,  Co3[Fe(CN)G]2  ,  insoluble  in  acids.     With  ?  mmo- 

nium    chloride    and    hydroxide,    excess    of   ferricyanide    gives    a    blood-  red 

solution,  a  distinction  of  cobalt,  from  nickel,  manganese  and  zinc. 
Copper,  a  yellow-green  precipitate,  Cu3[Fe(CN)0]2  ,  insoluble  in  HC1  . 
Gold,  no  precipitate. 

Iron  (ferrous),  dark  Hue  precipitate,  Fe3[Fe(CN)G]2  ,  insoluble  in  acids. 
Iron  (ferric),  no  precipitate,  a  darkening  of  the  liquid. 
Lead,  110  precipitate,  except  in  concentrated  solutions  (dark  brown). 
Manganese,  brown  precipitate,  Mn3[Fe(CN)ti]2  ,  insoluble  in  acids. 
Mercury  (mercurous),  red-brown  precipitate,  turning  w^hite  on  standing. 
Mercury  (mercuric),  no  precipitate. 
Nickel,  yellow-green  precipitate,  Ni3[Fe(CN)c]2  ,  insoluble  in  hydrochloric  acid. 

With   ammonium  chloride   and   hydroxide,   excess   of  ferricyanide   gives   a 

copper-red  precipitate. 

Silver,  a  red-brown  precipitate,  Ag)3Fe(CN),)  ,  soluble  in  NH4OH  . 
Tin  (stannous),  white  precipitate,  Sn3[Fe(CN)6J2  ,  soluble  in  hydrochloric  acid. 
Tin  (stannic),  no  precipitate. 
Uranium  (uranous),  no  precipitate. 
Zinc,  orange  precipitate,  Zn3[Fe(CN)6]2  ,  soluble  in  HC1  and  in  NH4OH  . 

Ferricyanides,  ferric  combinations,  are  capable  of  acting  as  oxidizing  agents^ 
becoming  ferrocyanides,  ferrous  combinations. 

4K3Fe(CN)6  +  2H2S  =  3K4Fe(CN)6  +  H4Fe(CN)6  +  S2 
2K3Fe(CN)6  +  2KI  =  2K4Fe(CN)6  +  I2  . 


Nitric  acid,  or  acidulated  nitrite,  by  continued  digestion  in  hot  solution, 
effects  a  still  higher  oxidation  of  ferricyanides,  with  the  production,  ;  moiig 
other  products,  of  nitroferricyanides  or  nitroprussides  (Play  fair,  Phil.  May.,  1845, 
(3),  26,  197,  271  and  348).  These  salts  are  generally  held  to  have  the  composi- 
tion represented  by  the  acid  H.,Fe(NO)  (CN)5  .  Sodium  iiitropmsside  is  u  ed  as 
a  reagent  for  soluble  sulphides  —  that  is,  in  presence  of  alkali  hydroxiies,  a 
test  for  hydrosulphuric  acid;  in  presence  of  hydrosulphuric  acid,  a  test  for 
alkali  hydroxides  (§207,  66). 

K3Fe(CN)6  is  reduced  to  K4Fe(CN)6  by  Pd  ,  Th  ,  Mg  and  As,  but  rot  by 
Pb  ,  Hg  ,  Ag-  ,  Sb  ,  Sn  ,  Au  ,  Pt  ,  Bi  ,  Cu  ,  Cd  ,  Te  ,  Al  ,  Fe  ,  Co  ,  Mn  ,  Zn  ai.d  In  . 
When  a  sheet  of  any  metal  except  Au  and  Pt  is  placed  in  contact  v  ith  a 
solution  of  K3Fe(CN)fi  and  FeCl3  ,  a  coating  of  Prussian  blue  is  soon  formed 
(Boettger,  J.  C.,  1873,  26,  473). 
Pb"  with  potassium  hydroxide  forms  PbO2  and  potassium  ferrocyanide  (Watts' 

Dictionary,  1889,  2,  340). 
Sn"  with  potassium  hydroxide  forms  potassium  stannate,  K2Sn03   and  uotas- 

sium  ferrocyanide  (Watts'  Dictionary,  I.  c.). 
Cr'"  forms  in  alkaline  mixture  a  chromate  and  a  ferrocyanide  (Bloxam,  C.  N., 

1885,  52,  109). 
Mn"    with    potassium    hydroxide    forms    Mn02    and    potassium    ferrocyanide 

(Boudault,  J.  pr.,  1845,  36,  23). 
Co"  and  Ni"  are  not  oxidized. 

In  alkaline  solutions  K3Fe(CN)6  oxidizes  sugar,  starch,  alcohol,  oxalic  acid 
and  indigo  (Wallace,  J.  C.,  1855,  7,  77;  Mercer,  Phil.  Mag.,  1847,  (3),  31,  1/G). 
HN02  and  HN03  both  form  hydronitroferricyanic  acid,  H2Fe(NO)  (CN)6  . 


§233.  CYANIC  ACID.  279 

NO  ii.  alkaline  solution  becomes  a  nitrate  (Wallace,  I.  r.). 

P  in  f  Ikaline  solution  becomes  a  phosphate  (Wallace,  /.  c.). 

HH2>'02  forms  H4Fe(CN)fi  and  H  PO,  . 

H2S  ^orms  S,  then  H..SO,  and  H,Fe(CN),;  (Wallace,  I.e.). 

SO,  forms  H2S04  and  H,Fe(CN)0  . 

Cl  decomposes  ferricyanidcs. 

HC1O,  acts  upon  K3Fe(CN)0  ,  forming-  potassium  superferricyanide,  K2Fe(CN)8 

(Skraup,  A.,  1877,  189,  :!(ks). 
HI  fcrms  H4Fe(CN)(J  and  I  . 

Ferricyanides  in  solution  are  detected  by  the  reactions  with  ferrous  and 
ferric  salts  (§126,  G7>).  Insoluble  compounds  are  ignited  (under  a  hood)  with 
a  fixe  1  alkali,  giving1  an  alkali  cyanide,  ferric  oxide,  and  an  oxide  of  the  metal 
in  co  nbination.  Detect  the  alkali  cyanide  as  directed  (§230,  8).  A  ferri- 
cyani  le  is  estimated  by  reduction  to  ferrocyanide  with  KI  in  presence  of  con- 
cent r.ited  HC1;  the  liberated  iodine  being  titrated  with  standard  Na2SaO3  . 
Or  it  is  reduced  to  ferrocyanide  by  boiling  with  KOH  and  FeS04  ,  filtering, 
acidulating  with  H2S04  and  titrating  with  KMnO4  . 


§233.  Cyanic  acid.     HCNO  =  23.018. 
H  —  0  —  C=N. 

The  cyanates  of  the  alkalis  and  of  the  fourth-group  metals  may  be  made  by 
passii  g  cyanogen  gas  into  the  hydroxides.  The  cyanates  of  the  alkalis  are 
easily  prepared  by  fusion  of  the  cyanide  with  some  easily  reducible  oxide. 

C2N2  -f  2KOH  =  KCNO  +  KCN  +  H20 

KCN  +  PbO  —  KCNO  +  Pb 

4KCN  +  Pb304  =  4KCNO  +  r,Pb 

The  free  acid  may  be  obtained  by  heating  cyanuric  acid,  H3C3N3O3  ,  to 
redness,  better  in  an  atmosphere  of  CO2  .  Cyanic  acid  is  found  in  the  dis- 
tillate H3C8N308  =  3HCNO  . 

Absolute  cyanic  acid,  HCNO  ,  is  a  colorless  liquid,  giving  off  pungent,  irri- 
tating vapor,  and  only  preserved  at  very  low  temperatures.  It  cannot  be 
forme  d  by  transposing  metallic  cyanates  with  the  stronger  acids  in  the  pres- 
ence >f  water,  by  which  it  is  changed  into  carbonic  anhydride  and  ammonia: 
HCNO  +  H2O  =  NH3  +  CO2  .  The  cyanates,  therefore,  when  treated  with 
hydrr-.'hloric  or  sulphuric  acid,  effervesce  with  the  escape  of  carbonic  anhydride 
(disti  -iction  from  cyanides),  the  pungent  odor  of  ci/iinic  acid  being  perceptible: 
2KCEO  +  2H_,SO4'  +  2H2O  =  K2S04  +  (NH4)2SO4  +  2C02 .  The  ammonia 
remains  in  the  liquid  as  ammonium  salt,  and  may  be  detected  by  addition  of 
potassium  hydroxide,  with  heat. 

The  cyanates  of  the  metals  of  the  alkalis  and  of  calcium  are  soluble  in  water; 
most  of  the  others  being  insoluble  or  sparingly  soluble.  All  the  solutions 
gradrtlly  decompose,  with  evolution  of  ammonia.  Silver  cyanate  is  sparingly 
soluble  in  hot  water,  readily  soluble  in  ammonia;  soluble,  with  decomposition, 
in  dibite  nitric  acid  (distinction  from  cyanide).  Copper  ci/anate  is  precipitated 
green  i  sh-y  ellow. 

Atni.innium  cyanate  in  solution  changes  gradually,  or  immediately  when  boiled, 
to  we- 1,  or  carbamide,  with  which  it  is  isomeric:  NH4CNO  =  CO(NH,),  .  The 
latter  is  recognized  by  the  characteristic  crystalline  laminae  of  its  nitrate, 
when  a  few  drops  of  the  solution,  on  glass,  are  treated  with  a  drop  of  nitric 
acid.  Also,  solution  of  urea  with  solution  of  mercuric  nitrate,  forms  a  white 
precipitate,  CH4N,O(HgO)2  ,  not  turned  yellow  (decomposed)  by  solution  of 
sodiuri  carbonate  (no  excess  of  mercuric  nitrate  being  taken).  Solution  of 
urea,  on  boiling,  is  resolved  into  ammonium  carbonate,  which  slowly  vapori/cs: 
CH4N,O  -f  2H2O  =  (NH.)  .CO;,  .  Cyanates,  in  the  dry  way,  are  reduced  by 
strong  deoxidizing  agents  to  cyanides. 

For  detection  of  a  cyanate  in  presence  of  cyanides,  see  Schneider,  B.,~  1895, 
28,  1540. 


280  THIOCJANIC  ACID.  §234. 

§234.  Thiocyanic  acid.     HCNS  =  59.088 . 
H  —  S  —  C  =  N. 

An  aqueous  solution  of  HCNS  may  be  obtained  by  treating1  lead  thiocyanate 
suspended  in  water  with  H2S  ,  also  by  treating  barium  thiocyanate  with  H2S04 
in  molecular  proportions.  The  anhydrous  acid  is  obtained  by  treating  dry 
Hg(CNS)2  with  H,S  .  Potassium  thiocyanate  is  formed  by  fusing  KCN  with 
S.  Or  two  parts  of  K4Fe(CN)e  with  one  part  of  sulphur.  Also  by  fusing  the 
cyanide  or  ferrocyanide  of  potassium  with  potassium  thiosulphate,  K2S203: 

2KCN  +  S2  =  2KCNS 

K4Fe(CN)G  +  3S2  =  4KCNS  +  Fe(CNS)2 

4KCN  +  4K2S203  =  4KCNS  -f  3K2S04  -f  K2S 

2K4Fe(CN)6  +  12K2S20S  =  12KCNS  +  9K2SO4  +  K2S  +  2FeS 

Thiocyanic  acid  is  quite  as  frequently  called  sulphocyanic  acid,  and  its  salts 
either  thiocyanates  or  sulphocyanates.  It  corresponds  to  cyanic  acid,  HCNO  , 
oxygen  being1  substituted  for  sulphur. 

Absolute  thiocyanic  acid,  HCNS  ,  is  a  colorless  liquid,  crystallizing  at  12° 
and  boiling  at  85°.  It  has  a  pungent,  acetous  odor,  and  reddens  litmus.  It  is 
soluble  in  water.  The  absolute  acid  decomposes  quite  rapidly  at  ordinary 
temperatures;  the  dilute  solution  slowly;  with  evolution  of  carbonic  anhydride, 
carbon  disulphide,  hydrosulphuric  acid,  hydrocyanic  acid,  ammonia,  and  other 
products. 

The  same  products  result,  in  greater  or  less  degree,  from  transposing  soluble 
thiocyanates  with  strong  acids;  in  greater  degree  as  the  acid  is  stronger  and 
heat  applied;  while  in  dilute  cold  solution,  the  most  of  the  thiocyanic  acid 
remains  undecomposed,  giving  the  acetous  odor.  The  thiocyanates,  insoluble 
in  water,  are  not  all  readily  transposed.  Thiocyanates  of  metals,  whose  sul- 
phides are  insoluble  in  certain  acids,  resist  the  action  of  the  same  acids. 

The  thiocyanates  of  the  metals  of  the  alkalis,  alkaline  earths;  also,  those  of 
iron  (ferrous  and  ferric),  manganese,  zinc,  cobalt  and  copper — are  soluble  in 
water.  Mercuric  thiocyanate,  sparingly  soluble;  potassium  mercuric  thiocyanate, 
more  soluble.  Silver  thiocyanate  is  insoluble  in  water,  insoluble  in  dilute  nitric 
acid,  slowly  soluble  in  ammonium  hydroxide. 

Solutions  of  metallic  thiocyanates  give,  with  soluble  salts  of: 
Cobalt,  very  concentrated,  a  blue  color,  Co(CNS)2  ,  crystallizable  in  blue 
needles,  soluble  in  alcohol,  not  in  carbon  disulphide.  The  coloration  is 
promoted  by  warming*,  and  the  test  is  best  made  in  an  evaporating  dish. 
In  strictly  neutral  solutions,  iron,  nickel,  zinc  and  manganese,  do  not 
interfere. 

Copper,  if  concentrated,  a  black  crystalline  precipitate,  Cu(CNS)2  ,  soluble  in 
thiocyanate.     With  sulphurous  acid,  a  white  precipitate,  CuCNS;  also  with 
hydrosulphuric    acid    (used    to    separate    a   thiocyanate    from    a    chloride) 
(Mann,  Z.,  1889,  28,  668). 
Iron  (ferrous),  no  precipitate  or  color. 

Iron  (ferric),  an  intensely  blood-red  solution  of  Fe(CNS)3  ,  decolored  by  solu- 
tion of  mercuric  chloride  (§126,  6&,  distinction  from  acetic  acid);  decolored 
by  phosphoric,  arsenic,  oxalic  and  iodic  acids,  etc.,  unless  with  excess  of 
ferric  salt;  decolored  by  alkalis  and  by  nitric  acid,  not  by  dilute  hydro- 
chloric acid.  On  introduction  of  metallic  zinc,  it  evolves  hydrosulphuric 
acid.  Ferric  thiocyanate  is  soluble  in  ether,  which  extracts  traces  of  it 
from  aqueous  mixtures,  rendering  its  color  much  more  evident  by  the 
concentration  in  the  ether  layer. 
Lead,  gradually,  a  yellowish  crystalline  precipitate,  Pb(CNS)2  ,  changed  by 

boiling  to  white  basic  salt. 

Mercury  (mercurous),  a  white  precipitate,  HgCNS  ,  resolved  by  boiling  into 
Hg  and  Hg(CNS)2  .  The  mercurous  thiocyanate,  HgCNS,  swells  greatly 
on  ignition  (being  used  in  "Pharaoh's  serpents"),  with  evolution  of  mer- 
cury, nitrogen,  thiocyanogen,  cyanogen  and  sulphur  dioxide. 


§235,  1.  NITROGEN.  281 

Mercury  (mercuric),  in  solutions  not  very  dilute,  a  white  precipitate, 
Hg(CNS)2  ,  somewhat  soluble  in  excess  of  the  thiocyanates,  sparingly 
soluble  in  water,  moderately  soluble  in  alcohol.  On  ignition,  it  swells  like 
the  mercurous  precipitate. 

Platinum.  1'latinic  chloride,  gradually  added  to  a  hot,  concentrated  solution 
of  potassium  thiocyanate,  forms  a  deep-red  solution  of  double  thiocyanate  of 
potassium  and  platinum  (KCNS),Pt(CNS)4  ,  or  more  properly,  K,Pt(CNS)r>, 
potassium  thfocyanoplatinate.  The  latter  salt  gives  bright-colored  precipi- 
tates with  metallic  salts.  The  thiocyanoplatinate  of  lead  (so  formed)  is 
golden-colored;  that  of  silver,  orange-red. 

Silver,  a  white  precipitate,  AgCNS  ,  insoluble  in  water,  insoluble  in  dilute 
nitric  acid,  slowly  soluble  in  ammonium  hydroxide,  readily  soluble  in  excess 
of  potassium  thiocyanate;  blackens  in  the  light;  soluble  in  hot  concentrated 
H,S04  (separation  from  Ag-Cl)  (Volhard,  A.,  1877,  190,  1). 

Certain  active  oxidizing  agents,  viz.,  nascent  chlorine,  and  nitric  acid  contain- 
ing nitrogen  oxides,  acting  in  hot,  concentrated  solution  of  thiocyanates,  pre- 
cipitate perthiQcyanogen,  H(CNS)3  ,  of  a  yellow-red  to  rose-red  color,  even  blue 
sometimes.  It  may  be  formed  in  the  test  for  iodine,  and  mistaken  for  that 
element,  in  starch  or  carbon  disulphide.  If  boiled  with  solution  of  potassium 
hydroxide,  it  forms  thiocyanate. 

Concentrated  hydrochloric  acid,  or  sulphuric  acid,  added  in  excess  to  water 
solution  of  thiocyanates,  causes  the  gradual  formation  of  a  yellow  precipitate, 
pcrtJiiori/aiiic  arid,  (HCN)2S3  ,  slightly  soluble  in  hot  water,  from  which  it 
crystallizes  in  yellow  needles.  It  dissolves  in  alcohol  and  in  ether. 

Potassium  thiocyanate  can  be  fused  in  closed  vessels,  without  decomposition; 
but  with  free  access  of  air,  it  is  resolved  into  sulphate  and  cyanate,  with 
evolution  of  sulphurous  acid. 

When  thiocyanic  acid  is  oxidized,  the  final  product,  as  far  as  the  sulphur  is 
concerned,  is  always  sulphuric  acid  or  a  sulphate.     In  many  cases  (in  acid  mix- 
ture)  it  has  been  proven  that  the   cyanogen  is  evolved  as  hydrocyanic  acid. 
In  other  eases  the  fame  reaction  is  assumed  as  probable. 
PbO2  and  Pb304  form  Pb"  and  sulphuric  acid,  in  acid  mixture  only  (Hardow, 

,/.  (7.,  1859,  11,  174). 

H3AsO4  forms  H3AsO3  ,  hydrocyanic  and  sulphuric  acids. 
Co'"  forms  Co"  ,  hydrocyanic  and  sulphuric  acids. 
Ni'"  forms  Ni"  ,  hydrocyanic  and  sulphuric  acids. 
Crvi  forms  Cr"'  ,  hydrocyanic  and  sulphuric  acids. 
Mn"+n  forms  Mn"  ,  hydrocyanic  and  sulphuric  acids.     In  alkaline  mixture,  a 

cyanate  and  sulphate  are  formed  (Wurtz's  Diet.  Chim.,  3,  95). 
HNOL)  forms  sulphuric  acid  and  nitric  oxide. 
HNO3  forms  sulphuric  acid  and  nitric  oxide. 

Cl  forms  at  first  a  red  compound  of  unknown  composition,  then  HC1 ,  H2S04 
and  HCN  are  produced.  In  alkaline  mixture  a  chloride  and  sulphate  are 
formed. 

HC10  same  as  with  Cl  . 

HC1O3  forms  sulphuric,  hydrochloric  and  hydrocyanic  acids. 
Br  forms  HBr  and  H2S04;  but  with  alkalis,  a  bromide  and  sulphate. 
HBrO8  forms  HBr  and  H2SO4  . 
HIO3  forms  H2S04  and  free  iodine. 


§235.  Nitrogen.     N  =  14.01.     Valence  one  to  five  (§11). 

1.  Properties. — Weight  of  molecule,  N2  ,  28.08.  Vapor  density,  14  (Jolly,  TF. 
A.,  1879,  6,  536).  At  — 123.8°,  under  pressure  of  42.1  atmospheres,  it  condenses 
to  a  liquid  (Sarrau,  C.  r.,  1882,  94,  718).  Boiling  point,  —194.4°  (Olszewski,  W.  A., 
1897,  31,  58).  Liquid  nitrogen  is  colorless  and  transparent.  The  gas  is  taste- 
less, odorless  and  colorless.  Not  poisonous,  but  kills  by  excluding  air  from  the 
lungs.  Does  not  burn  or  support  combusion.  It  is  very  inert,  not  attacking 
other  free  elements.  Its  simplest  combinations  are  the  following:  N— '"H'3  , 
N2O  ,  NO  .  N2O3  ,  NO.,  and  N,O,  .  The  number  of  organic  compounds  contain- 
ing nitrogen  is  very  large.  The  nitrogen  in  all  compounds  that  are  the 


282  ^HYDRAZOIC  ACID.  §235,2. 

immediate  products  of  vegetable  growth  bns  a  valence  of  minus  three  and 
may  without  change  of  bonds  be  converted  into  N— '"H'a  .  This  statement  is 
made  with  a  limited  knowledge  of  the  facts  and  without,  at  present,  having 
conclusive  proof;  and  merely  predicting  that  future  research  will  verify  it. 

2.  Occurrence. — It  constitutes  about  four-fifths  of  the  volume  of  the  atmos- 
phere.    It  occurs  as  a  nitrate  in  various  salts  and  in  various  forms  as  a  con- 
stituent of  animal  and  vegetable  growths. 

3.  Formation. — (a)    From    the    air,    the    oxygen    being1   removed    by    red-hot 
copper,  the  C02  by  potassium  hydroxide,  the  ammonia  and  water  by  passing 
through  H2SO4.*     (6)  Ignition  of    ammonium    dichromate,    *NH4)2Cr2O7  =  N2  + 

Cr2O3  +  4H2O  .  (c)  By  heating  ammonium  nitrate  and  peroxide  of  manganese 
to  about  200°  (Gatehouse,  C.  N.,  1877,  36,  118).  (d)  Ignition  of  NH4C1  and 
K2Cr2O7:  2NH4C1  +  K2Cr2O7  =  ^KCl  +  N2  +  Cr2O3  +  4H2U  .  Unless  the  tem- 
perature be  carefully  guarded,  traces  of  NO  are  formed,  which  may  be  removed 
by  passing  the  gases  through  FeSO4 .  (e)  Action  of  chlorine  upon  NH3:  8NH.3 
-j-  OC12  =  6NH4C1  +  N2 .  The  NH3  must  be  kept  in  excess  to  avoid  the  forma- 
tion of  the  dangerously  explosive  chloride  of  nitrogen,  NC13  .  (/)  Removing 
the  oxygen  from  the  air  by  shaking  with  NH4OH  and  copper  turnings.*  (g) 
Burning  phosphorus  in  air  over  water.*  (h)  By  passing  air  through  a  mixture 
of  FeS  and  sawdust;  then  through  a  pyrogallate  solution,  and  finally  through 
concentrated  H2SO4.*  (i)  By  shaking  air  with  Fe(OH)2  and  Mn(OH)2.* 
(j)  By  passing  air  through  an  alkaline  pyrogallate.*  (/c)  By  passing  air,  from  which 
UOohas  been  removed,  mixed  with  hydrogen  over  heated  platinum  black,  the  hydro- 
gen having  been  added  in  just  sufficient  quantity  to  form  water  with  all  the  oxygen  * 
(Damoulin,  «/.,  1851,  321).  (Z)  By  warming  a  concentrated  solution  of  NH4NO2  or 
a  mixture  of  KNO2  and  NH4C1:  NH4NO2  =  N2  +  2H2O  .  Potassium  dichromate  is 
added  to  oxidize  to  nitric  acid  any  of  the  oxides  of  nitrogen  that  may  be  formed 
(Gibbs,  B.,  1877, 1387).  (ra)  By  action  of  potassium  or  sodium  hypobromite  upon 
ammonium  chloride:  3NaBrO  +  2NH4C1  =  N2  +3NaBr  +  l^HCl  -f  3H2O  . 

4.  Preparation. — Nitrogen  has  been  economically  produced  by  most  of   the 
above  methods. 

5.  Solubilities. — Nitrogen  is  nearly  insoluble  in  all  known  liquids. 

6.  Reactions. — At  ordinary  temperatures  nitrogen  is  not  acted  upon  by  other 
compounds.  Nodules  containing  the  so-called  nitrifying  bacteria  growing  on  the  roots 
of  leguminous  plants  absorb  nitrogen  and  build  up  nitrogenous  compounds  therewith. 

7.  Ignition. — Under    electric    influence    it    combines    slowly   with   hydrogen; 
also  with  B  ,  Cr  ,  Mg  ,  Si  and  V  . 

8.  Detection. — Nitrogen  is   more  easily  detected  by  the  nature  of  its  com- 
pounds than  by  the  properties  of  the  liberated  element. 

9.  Estimation.  — (a)  As  free   nitrogen  by  measuring  the  volume  of  the  gas. 
(6)   By  oxidation  of  the  organic  substance  with  hot  concentrated  H2SO4  ,  which 
also   converts   the   nitrogen   into   ammonium   sulphate.     For   details   see   works 
on  organic  analysis,      (c)  By  decomposition  of  the  organic  material  with  potas- 
sium   permanganate    in    strong    alkaline    solution,    forming    ammonia,     (d)  By 
combustion  of  the  organic  compound  in  presence  of  CuO  and  Cu°.    absorbing 
the  CO2  by  KOH  and  determining  the  nitrogen  by  volume. 

(For  Hydroxylamine,  see  foot-note,  page  286.) 

§236.  Hydrazoic  acid  (Azoimide).     N3H  =  43.038. 

N\ 
Constitution,  II    >NH 

.  W 

Curtius,  B.,  1890,  23,  3023.  A  clear  mobile  liquid  of  penetrating  odor,  a 
very  irritative  effect  upon  the  nostrils  and  the  skin,  and  readily  exploding 
with  exceeding  violence.  Boiling  point,  about  37°.  Soluble  in  water  ;  nd 
alcohol.  An  acid  of  marked  acidity,  dissolving  a  number  of  metals  with 
*  Nitrogen  made  from  the  air  is  not  pure.  It  contains  about  one  per  cent  of  argon  and 
smaller  amounts  of  krypton,  neon,  and  xenon.  Because  of  the  presence  of  these  impurities 
its  density  is  greater  than  that  of  nitrogen  prepared  from  chemical  compounds.  (Ramsey 
and  Rayleigh.) 


£238,  <;.  NITROUS    OX1DK— \1TRIC   OXIDE. 

evolution  of  hydrogen.  Its  salts,  the  trinitrides  of  the  metals  of  the  alkalis 
and  the  alkaline  earths,  are  soluble  in  water  and  crystallizable  (Dennis,  J.  Am. 
/S'oc.,  1898,  20,  225).  Potassium  trinitride  precipitates  from  thorium  salts,  the 
hydroxide  of  this  metal  in  quantitative  separation  from  cerium,  lanthanum, 
neodymium  and  praseodymium  (Dennis,  J.  Am.  Nor..  1S9<>,  18,  947).  Hydro- 
nitric  acid  is  formed  by  treating1  ammonia  with  sodium,  and  the  resulting1 
sodamide,  NaNHo  ,  with  nitrous  oxide:  2NaNH.,  +  N20  =  NaN3  +  NaOH  -j- 
NH3  (Wislicenus,  J3.,  1892,  25,  2084). 

§237.  Nitrous  oxide.     N20  =  44.02 . 
!T20-",  N  —  0  —  N  . 

Nitrous  oxide  becomes  a  colorless  liquid  at  0°  under  pressure  of  three 
atmospheres  (Farady,  A.,  1845,  56,  157).  Melts  at  —99°  and  boils  at  —92° 
(\\ 'ills,  -/.  C.,  1S74,  27,  21).  It  is  a  colorless  gas  with  slight  sweetish  smell  and 
taste.  Supports  combustion.  When  breathed  acts  as  an  anaesthetic  of  short 
duration;  and  is  used  in  dentistry  for  that  purpose.  Decomposed  by  heat 
completely  at  900°  into  N  and  O  (Meyer,  PyrocJiemisch.  Untersuch.,  1885).  Passed 
over  red-hot  iron  N  and  Fe^Os  are  formed.  K  and  Na  burn  in  nitrous  oxide, 
liberating  the  nitrogen.  As  a  rule  both  gases  and  solids  that  burn  in  air  burn 
also  in  nitrous  oxide.  It  is  formed:  (a)  By  heating  ammonium  nitrate  in  a 
retort  from  170°  to  260°:  NH4NO,  =  N2O  +  2H,O  .  (ft)  By  passing  NO  through 
solution  of  S02  .  (c)  By  action  of  HNO3;  sp.  (jr.,  1.42,  diluted  with  an  equal 
volume  of  water,  upon  metallic  zinc,  (d)  A  mixture  of  five  parts  of  SnCl2  ,  ten 
parts  of  HC1 ,  sp.  (jr.,  1.21,  and  nine  parts  of  HNO3  ,  sp.  (jr.,  1.3,  is  heated  to 
boiling:  2HN03  +  4SnCl2  +  8HC1  =  4SnCl4  +  N,O  +  5H2O  (Campari,  J.  C., 
1889,  55,  569). 

§238.  Nitric  oxide.    NO  =  30.01 . 
N"0-",  N  =  0  . 

1.  Properties. — The  vapor  density  (15)  shows  the  molecule  to  be  NO  (Daccomo 
and    Meyer,    B.,    1887,    20,    1832).     Under    pressure    of    one    atmosphere    it    is 
liquified  at  — 153.6°,   and   under  71.2   atmospheres   at  — 93.5°,   and  solidifies   at 
—167°  (Olszewski,  C.  r.,  1877,  85,  1016).  Odor  and  taste  unknown,  on  account  of 
its  immediate  conversion  into  NO2  on  exposure  to  the  air. 

2.  Occurrence. — Not  found  free  in  nature. 

3.  Formation. — (</)    Reduction  of  nitric  acid  by  means  of   ferrous  sulphate 
previously  acidulated  with  H2SO4  .     (ft)   Action  of  cold  nitric  acid,  sp.  gr.,  1.2, 
upon  metallic  copper;  unless  great  care  be  used  other  oxides  of  nitrogen  are 
produced,     (c)    SO2    is   passed  into   slightly   warmed   HNO3  ,   sp.   gr.,   1.15,   and 
excess  of  SO2    removed  by  passing  through  water,     (d)    According  to   Emich 
(.I/.,   1893,   18,  73),  a  strictly   pure  nitric  oxide  is  made  by  treating  mercury 
with  a  mixture  of  nitric  and  sulphuric  acids. 

4.  Preparation. — Same  as  above.  i 

5.  Solubilities. — Soluble  in  about  ten  volumes  of  water  and  in  five  volumes 
of  nitric  acid,  sp.  gr.,   1.3.     One  hundred  volumes  of  H2S04  ,  sp.  gr.,  1.84,   and 
1.50,  dissolve  3.5  and  1.7  volumes  respectively    (Lunge,  B.,   1885,   18,   1391).     A 
16  per  cent  solution  of  ferrous  sulphate  dissolves  six  times  its  own  volume  of 
the  gas  forming  the  "  brown  ring,"  which  is  decomposed  at  100°.     Soluble  in 
CS2  and  in  alcohol. 

6.  Reactions. — When  heated  in  nitric  oxide  to  450°,  Ag  ,  Hg-  and  Al  are  un- 
changed; filings  of  Cu  ,  Fe  ,  Cd  and  Zn  are  superficially  oxidized,  but  lead  is 
completely  changed   to  PbO;  while  if   the  metals  are  in   an    exceedingly   line 
state  of  division   (by  reduction  of  their  oxides  by  hydrogen),  Ni  at  200°   be- 
comes NiO  ,  Fe  at  200°  forms  FeO  ,  Cu  at  200°  forms  Cu20:  the  higher  oxides  of 
these  metals  not  being  thus  produced  (Sabatier  and  Senderens,  C.  r.,  1892,  114, 
1429).     Oxidized  to  KN03  by  KMnO4:  KMnO,  +  NO  =  MnO,  +  KNO3   (Wank- 
lyn  and  Cooper,  Phil.  Mag.,\S18,  (5),  6,  288). 


284  NITROUS  ACID.  §23&,  1. 

§239.  Nitrous  acid.     HN02  =47.018. 
H'N'"0-"2,  H  —  0  —  N  =  0. 

1.  Properties. — Nitrous    acid    is   known    only    in    solution.     Made    by    adding 
N2O3  to  water.     It  has  a  blue  color  and,  owing  to  its  tendency  to  dissociation 
(6HNO2  =  2HN03  +  4NO  +  2H20),  is  very  ^stable  (Fremy,  C.  r.,  1870,  70,  61). 
Nitrous  anhydride  is  obtained  when  a  mixture  of  one  volume  of  oxygen  and 
four  volumes  of  nitric  oxide  are  passed  through  a  hot  tube,  4NO  -f-  02  —  2N203 . 
It  is  a  deep  red  gas,  condensing  to  a  blue  liquid  at  14.4°  under  755  mm.  pressure 
(Gains,  C.  N.,  1883,  48,  97). 

2.  Occurrence. — Traces  of  ammonium  nitrite   are  found   in   the   air,   in   rain 
\vater,  river  water  and  in  Chili  saltpeter.     When  found  in  nature  it  is  usually 
accompanied  by  nitrates. 

3.  Formation. — By  action  of  nitric  acid,  sp.  gr.,  1.35,  upon  starch  or  arsenous 
oxide.     At   70°   nearly  pure   N203    is  obtained,   which   passed   into   cold   water 
forms  HNO.,  .     Nitrites  of  potassium  and  sodium  may  be  formed  by  ignition 
of  their  nitrates   (a  prolonged  high  heat  forming  the  oxides).     Or  the   alkali 
nitrites   may   be  made   by   fusing  the   nitrates   with   finely   divided   iron;   lead 
nitrite  by  fusing  lead  nitrate  with  metallic  lead,    and   silver  nitrite   may   be 
made  from  these  by  precipitation;  and  from  this  salt  many  nitrites  may  be 
made  nearly  pure  by  transposition;    e.g.,   Bad,    +    2AgNO*.    =  Ba(NO.>)«    + 
2AgCl  and  then  Ba(NO2)2  +  ZnSO,  =  Zn(N02)2  +  BaSO4  . 

4.  Preparation. — Same  as  above. 

5.  Solubilities. — Silver   nitrite   is   only   sparingly   soluble    (120   parts   of   cold 
water).     The  other  normal  nitrites  are   soluble;   but  many  basic  nitrites   are 
insoluble. 

Nascent  hydrogen  in  presence  of  an  alkali  reduces  nitrates  to  nitrites;  e.  g., 
sodium  amalgam,  aluminum  wire  in  hot  KOH  ,  etc.  Used  in  excess  the  nascent 
hydrogen  reduces  the  nitrogen  still  further,  forming  NH3  . 

G.  Reactions. — A. — With  metals  and  their  compounds. — Nitrous  acid  acts 
sometimes  as  an  oxidizer,  sometimes  as  a  reducer;  in  the  former  case  NO  is 
iiHUttlly  produced  (under  some  conditions  N2O  ,  N  and  NH3  are  formed);  in  the 
latter  case  nitric  acid  is  the  usual  product,  but  sometimes  NO2  is  produced. 

1.  Pb02  becomes  Pb"  and  nitric  acid. 

2.  Hg'  becomes  Hg°  and  nitric  acid. 

3.  Crvi  becomes  Cr'"  and  nitric  acid. 

4.  Co"  becomes  Co'"  and  nitric  oxide.     Excess  of  KN02   with  acetic  acid  is 
used  to  separate  cobalt  from  nickel  (§132,  (k?). 

5.  Ni"'  becomes  Ni"  and  nitric  acid. 

6.  Mn"  +  n  becomes  Mn"  and  nitric  acid. 

B. — With  non-metals  and  their  compounds. — 

1.  H4Fe(CN)6  becomes  first  H3Fe(CN)«  and  then  hydronitroferricyanic  acid. 
Solution  of  indigo  in  sulphuric  acid  is  bleached  by  nitrites. 

2.  Nitrites  are  decomposed  by  nitric  acid. 

3.  HH,PO,  becomes  H3PO4  and  NO. 

J/.  H,S  does  not  displace  or  transpose  alkali  nitrites,  but  if  acetic  acid  be 
added  to  liberate  the  nitrous  acid,  then  S°  and  NO  are  produced.  H2S03  be- 
comes ELSO4  and  chiefly  NO  .  With  excess  of  H,SO3  ,  N2O  or  NH3  is  formed. 
See  Weber,  For///.,  I860,  127,  543,  and  1867,  130,  277;  Fremy,  C.  r.,  1870,  70,  61. 

5.  HC10:1  becomes  Cl°  and  HN03  . 

G.  HBr03  becomes  Br°  and  HNO3 

7.  HI  becomes  1°  and  NO  . 
HIO3  becomes  1°  and  HNO3  . 

7.  Ignition.— In  general  nitrites  are  changed  to  oxides,  but  with  potassium 
and  sodium  nitrites  a  white  heat  is  required,  and  with  nitrites  of  Ag .  Hg  , 
Au  and  Pt  the  dissociation  goes  a  step  further,  the  free  metals  being  produced. 

8.  Detection.  —  (1)   Formation   of   brown   ring   when   ferrous  sulphate  solution 
and  a  nitrite  is  acidulated  with  acetic  acid.     Nitrates  require  a  stronger  acid  for 


§241,  1.  NITROGEN    PEROXIDE — NITRIC    ACID.  285 

their  transposition.  (2}  A  mixture  of  a  nitrite  and  KI  liberates  iodine  on  adc'i- 
lion  of  acetic  acid  (nitrates  requiring  a  stronger  acid  for  transposition).  (3) 
Nitrous  acid  with  iodic  acid  liberates  iodine,  and  nitric  acid  is  produced.  (4) 
Solution  of  potassium  permanganate  acidified  with  sulphuric  acid  is  reduced 
by  nitrites  (distinction  from  nitrates). 

9.  Estimation. — Acidify  with  acetic  acid,  distil  and  titrate  the  distillate  with 
standard  solution  of  permanganates, 

§240.  Nitrogen  peroxide  (dioxide).    N02  =  46.01 . 

Vapor  density,  2.'!  (  R-misay,  ./.  ('..  1S!)0,  57,  f>00).  Melting-  point,  — 10C 
(Deville  and  Troost,  C.  r.,  1867,  64,  :.>:>?).  Boils  at  :>l.r>4°  (Thorpe,  J.  C.,  18b(). 
37,  224).  Below  10°  it  is  a  white  crystalline  solid.  Between  —10°  and  21.fi  ° 
a  liquid;  nearly  colcrless  at  !) °,  yellow  at  0°.  At  21.04°,  orange,  growing 
nearly  black  as  1h?  temperature  rises.  The  gas  does  not  support  combust io  , 
of  ordinary  fuels,  and  is  poisonous  when  inhaled.  It  dissolves  in  water,  form- 
ing a  greenish-blue  solution  containing  nitrous  and  nitric  acids.  With  an 
aqueous  solution  of  a  fixed  alkali  a  nitrate  and  nitrite  are  formed:  2N02  + 
2KOH  =  KN02  +  KNO3  +  H20  . 


§241,  Nitric  acid.     HN03  =  63.018. 

0 

II 
H'NvO-"3 ,  H  —  0  —  N  =  0  . 

1.  Properties. — Nitric   anhydride,  N20-  ,   is   a  colorless   solid,   melting  at   C.0° 
with  partial   decomposition   to   N02    and   O,   and   if   exposed   to   direct   sunlight 
decomposition  begins  at  lower  temperatures. 

Nitric  acid,  HN03  ,  has  not  been  perfectly  isolated;  that  containing  9D.8  per 
cent  of  HN03  is  a  colorless  highly  corrosive  liquid  (Roscoe,  A.,  1860,  116,  211), 
solidifies  at  47°  (Berthelot),  boils  at  86°,  but  dissociation  begins  at  a  lower 
temperature  and  is  complete  at  255°:  4HNO3  =  4N02  +  2!L,O  -f-  02  (Carius, 
B.,  1871,  4,  82S).  If  the  very  dilute  acid  be  boiled,  it  becomes  stronger,  and 
if  a  very  strong  acid  be  boiled  it  becomes  weaker,  in  both  cases  a  sp.  (jr.  of 
1.42  and  boiling  point  of  120°  being  reached;  the  acid  then  contains  about  70  per 
cent  of  HNO3  (Kolbe,  A.  Oh.,  1867  (4),  10,  136).  This  is  the  acid  usually 
placed  on  the  market.  The  reagent  usually  employed  has  a  sp.  (;r.  of  1.2 
(Fresenius  standard).  The  so-called  fuming  acid  has  a  specific  gravity  of  1.50 
to  1.52.  The  stronger  acid  should  be  kept  in  a  cool  dark  place  to  avoid  decom- 
position. 

2.  Occurrence. — Found  in  nature  as  nitrates  of  K,  Na ,  NH4  ,  Ca  ,  Mg  ,  and 
of   a   few   other   metals,    the    most    abundant   supply   coming   from    Chili    and 
Bolivia  as  sodium  nitrate,  "  Chili  saltpeter." 

3.  Formation.— (a)    Oxidation    of    nitrogenous    matter    in    presence    of    air, 
moisture  and  an  oxide  or  alkali;    (&)    by  oxidation   of   NO,   N203    or  NO,    by 
oxygen  (or  air)  in  presence  of  moisture;   (c)  from  NH3  ,  by  passing  a  mixture 
of  NH3  and  oxygen  through  red-hot  tubes. 

4.  Preparation. — By  treating  nitrates  with  sulphuric  acid  and  distilling. 

Nitrates  may  be  made:  (a)  By  dissolving  the  metal  in  nitric  acid,  except 
those  whose  metals  are  not  attacked  by  that  acid,  e.  g.,  An  ,  Pt ,  Al  and  Cr  ; 
and  also,  antimony  forms  Sb205 ,  arsenic,  H3As04  and  with  excess  of  hot 
acid  tin  forms  metastannic  acid  H10Sn5015  .  (&)  By  adding  HNO.,  to  the 
oxides,  hydroxides  or  carbonates.  All  the  known  nitrates  can  be  made 


286  CITRIC  ACID.  §241,  5. 

in  this  manner,  (c)  By  long  continued  boiling,  the  chlorides  of  all  ordi- 
nary metals  are  completely  decomposed,  no  chlorine  remaining,  except 
the  chlorides  of  Hg ,  Ag  ,  Au  and  Pt ,  which  are  not  attacked,  and  the 
chlorides  of  tin  and  antimony,  which  are  changed  to  oxides.  (Wurtz, 
Am.  S.,  1858,  75,  371;  Johnson,  Proc.  Am.  Ass.  ScL,  1894,  163.) 

The  anhydride  is  made:  (a)  By  passing  chlorine  over  silver  nitrate: 
4AgN03  +  2C12  ==  4AgCl  +  2N205  +  02 .  (&)  By  adding  anhydrous  P205 
to  HN03:  2NH03  +  P205  ==  2HP03  +  N205 . 

5.  Solubilities. — All  normal  nitrates  are  soluble.     A  few  are  decom- 
posed by  water,  e.  g.,  Bi(N03)3   +  H20  =  =  BiON03   +   2HNO;,  .     Most 
nitrates  are  less  soluble  in  nitric  acid  than  in  water,  e.  g.,  Cd  ,  Pb  ,  Ba  ,  etc.; 
the  barium  nitrate  being  completely  insoluble  in  HN03 ,  sp.  gr.,  1.42. 

Nitric  acid  decomposes  the  sulphides  of  all  ordinary  metals,  except 
mercuric  sulphide  which  "by  long  continued  boiling  with  the  concentrated 
acid  becomes  2HgS.Hg(N03)2 ,  insoluble  in  the  acid. 

6.  Reactions.     .4. — With  metals  and  their  compounds. — Nitric  acid  is 
a  powerful  oxidizer  but  unless  warmed  acts  more  slowly  than  chlorine. 
It   can  never   be   a   reducer.     The   following   products   are   formed:    H, 
NH8 ,  H2NOH  *,  N ,  N20  ,  NO  ,  HN02 ,  N02 .    If  the  acid  is  concentrated, 
in  excess  and  hot,  the  product  is  usually  entirely  nitric  oxide,  colorless, 
but  changing  to  the  red  colored  N02  by  coming  in  contact  with  the  air. 
Excess  of  the  reducer,  low  temperatures  and  dilute  solutions  favor  the 
production  of  nitrogen  compounds  having  lower  valence  and  of  hydrogen. 
Xascent  hydrogen  usually  forms  NH:i ,  always  the  ultimate  product  if  the 
hydrogen  be  produced  in  alkaline  mixture. 

Nitric  acid  oxidizes  all  ordinary  metals.  (It  does  not  act  upon  chro- 
mium, gold  or  platinum.)  It  forms  nitrates,  except  in  the  case  of  tin, 
antimony,  and  arsenic,  with  which  it  forms  H10Sn5015 ,  Sb20ri ,  and  H3As04  . 
With  the  respective  metals  it  forms  Hg'  or  Hg",  Sn"  or  Sn"",  As'"  or  Asv, 
Sb'"  or  Sbv,  Fe"  or  Fe'",  according  to  the  amount  of  nitric  acid  employed. 
With  copper  it  forms  cupric  nitrate  (never  cuprous);  with  cobalt  it  forms 
cobaltous  nitrate. 

*  Hydroxylamine,  NH2OU,  is  formed  by  the  reducing  action  of  Sn  and  HC1  upon  NO,N2O3, 
HXO3,  etc.  (Lessen,  A.,  1888, 252, 170);  also  by  the  action  of  H2S,  SO2,  K,  ]Vn,  M  g,  Zii,  and  Al  upon 
HXO3.  or  by  the  action  of  H2S  upon  certain  nitrates  (Divers  and  Haga,  C.  jV.,  1886,  54,  271 1.  By 
action  of  sodium  amalgam  upon  s  jdium  nitrite  solution,  XH2OH  is  produced  along  with  nitrous 
oxide,  free  nitrogen,  ammonia,  sodium  hyponi trite,  and  sodium  hydroxide,  the  highest  yield  of 
the  hydroxylamine  being  obtained  when  the  nitrite  solution  is  as  dilute  as  one  in  fifty,  the  mix- 
ture being  kept  cold  (Divers,  J.  C.,  1899,75,  87  and  89).  It  is  a  base  with  an  alkaline  reaction 
and  is  a  strong  reducing  agent.  When  pure  it  is  a  crystalline  solid,  odorless,  melting  at  33.05°, 
boiling  at  58°  at  22  mm.  pressure;  oxidized  by  oxygen  to  HNOa  (Lobry  de  Bruyn,  B.,  1892,  23; 
3,  190  and  684).  It  is  a  good  antiseptic  and  preservative.  It  combines  with  acids  to  form  salts: 
NH2OH  +HC1  =NH2OH.HC1 .  Hydroxylamine  hydrochloride  is  decomposed  by  alkalis  form- 
ing the  free  base,  which  is  decomposed  by  the  halogens,  KMnO4  ,  K2Cr2O?  ,  BaCh  and  PbO? 
Its  solution  in  ether  reacts  with  sodium,  forming  a  white  precipitate  ofNH?ONa, 


§241,  B,  8.  'NITRIC  ACID.  287 

1.  Pb02  is  not  changed.     Pb304  is  changed  thus:    Pb304  +  4HNO,  = 
PbO,  +  2Pb(NO ..),,  +  2K.O  . 

2.  Hg'  becomes  Hg". 

3.  Sn"  becomes  Sniv.     Stannous  chloride  and  hydrochloric  acid,  heated 
with  a  nitrate,  form  stannic  chloride,  and  convert  nitric  acid  to  ammonia 
(which  remains  as  ammonium  salt).     See  §71,  6c. 

4.  Sb'"  becomes  Sbv,  forming  Sb20.  ,  insoluble. 

5.  As'"  becomes  Asv,  forming  H3As04 . 

6.  Cu'  becomes  Cu". 

7.  Fe"  becomes  Fe'". 

B. — With  non-metals  and  their  compounds. 

1.  Carbon  (ordinary,  not  graphite)  becomes  C02  if  the  nitric  acid  be 
hot  and  concentrated. 

H2C204  becomes  C02 ,  in  hot  concentrated  acid. 

H4Fe(CN)G  becomes  first  H3Fe(CN)6  and  then  hydronitroferricyanic  acid. 

HCNS  is  oxidized,  the  sulphur  becoming  H2S04 . 

2.  Nitrites  are  all  decomposed,  nitrates  being  formed,  the  nitric  acid 
not  being  reduced.     The  nitrous  acid  liberated  immediately  dissociates: 
3HN02  =  2NO  +  HN03  -f-  H20  . 

S.  P°,  PH:; ,  HH2P02  and  HgPOg  become  H3P04  .  That  is  Pv~»  becomes 
Pv. 

4.  S  becomes  H,S04. 

H2S  becomes  first  S°  and  then  H,S04  . 
H,SO.{  becomes  H,S04    ;  and  in  general*SVI-n  becomes  SVI. 
2HNO3  +  3H2SO3  =  2NO  +  3H2SO4  +  H2O . 

5.  HC1,   nitrohydrochloric  acid:   2HN03  +  6HC1  =  2NO  +  4H20  +  3C12 
(Koninck  and  Nihoul,  Z.  anorg.,  1890,  477).     See  §269,  6#2. 

HC103  is  not  reduced.  Chlorates  are  all  transposed  but  not  decom- 
posed until  the  temperature  and  degree  of  concentration  is  reached  that 
would  dissociate  the  HC103  if  the  nitric  acid  were  absent. 

G.  Br°  is  not  oxidized.  HBr  becomes  Br°  and  is  not  further  oxidized. 
All  bromates  are  transposed  but  the  HBr03  is  not  decomposed  until  a  tem- 
perature and  degree  of  goncentratioii  is  reached  that  would  cause  the 
dissociation  of  the  HBr03  if  the  nitric  acid  were  absent. 

7.  1°  becomes  HI03 ,  very  slowly  unless  the  fuming  acid  be  used. 
HI  becomes  first  1°  ;  then  as  above. 

2HNO3  +  GHI ,  excess  =  2NO  +  3I2  +  H2O 
HI  +  HNO3 ,  excess  =  2NO  +  HIO3  +  H2O 

8.  In  general,  organic  compounds  are  oxidized.     Straw,  hay,  cotton,  etc., 
are  inflamed  by  the  strong  acid   (Kraut,  B.,  1881,  14,  301).     For  action 
on  starch,  see  Lunge,  B.,  1878,  11,  1229,  1641.    With  many  oi-ganic  bodies 


288  NITRIC  ACID.  §241,  7. 

substitution  products  are  formed,  the  oxides  of  nitrogen  taking  the  place 
of  the  hydrogen. 

7.  Ignition.— Nitric  acid  is  dissociated  by  heat:   4HNO,  =  4NO,  +  2H,O  +  Oo , 
complete  if  at  256°  (Carius,  B.,  1871,  4,  828).      No  nitrates  are  volatile  as  such. 
Ammonium  nitrate  is  dissociated:   NH4NO3  =   N2O  +  2H2O.   Some  nitrates,  e.  g., 
those  of  K  and  Na. ,  are  first  changed  to  nitrites  with  evolution  of  oxygen  onlv, 
and  at  an  intense  white  heat  further  changed  to  oxides  with  evolution  of  N.,*6 
as  well  as  oxygen.     As  a  final  result  of  ignition  the  nitrates  of  all  ordinary 
metals  are  left  as  oxides,  except  that  those  of  Hg  ,  A.g  ,  Au  and  Pt  are  reduced 
to  the  free  metal. 

A  mixture  of  potassium  nitrate  and  sodium  carbonate  in  a  state  of  fusion 
is  a  powerful  oxidizer;  e.g.,  changing  Sn"  to  Sniv  ,  As"'  to  Asv  f  gb'"  to  Sbv  , 
Fe"  to  Fe'"  ,  Cr'"  to  Crvi ,  Mnvi-n  to  Mnvi ,  svi-n  to  Svi ,  etc. 

Heated  on  charcoal,  or  with  potassium  cyanide,  or  sugar,  sulphur  or  other 
easily  oxidizable  substance  (as  in  gunpowder),  nitrates  are  reduced  with 
deflagration  or  explosion,  more  or  less  violent.  With  potassium  cyanide,  on 
platinum  foil,  the  deflagration  is  especially  vivid.  In  this  reaction  free  nitrogen 
is  evolved. 

Strongly  heated  with  excess  of  potassium  hydroxide  and  sugar  or  other 
carbonaceous  compound,  in  a  dry  mixture,  nitrates  are  reduced  to  ammonia, 
which  is  evolved,  and  may  be  detected.  In  this  carbonaceous  mixture,  the 
nitrogen  of  nitrates  reacts  with  alkalis,  like  the  unoxidized  nitrogen  in  car- 
bonaceous compounds. 

8.  Detection.  — Most  of  the  tests  for  the  identification  of  nitric  acid  are 
made  by  its  deoxidation,  disengaging  a  lower  oxide  of  nitrogen,  or  even, 
by  complete  deoxidation,  forming  ammonia. 

If,  with  concentrated  sulphuric  acid,  a  bit  of  copper  turning,  or  a  crystal 
of  ferrous  sulphate,  is  added  to  a  concentrated  solution  or  residue  of 
nitrate,  the  mixture  gives  off  abundant  brown  vapors ;  the  colorless  u  itric 
oxide,  NO ,  which  is  set  free  from  the  mixture,  oxidizing  immediately  in 
the  air  to  nitrogen  peroxide,  N02  : 

2KN03  -f  4H.SO,  +  3Cu  =  K2S04  +  3CuSO4  +  4H20  +  2NO 
2KN03  +  4H,S04  +  GFeS04  =  K2SO4  +  3Fea(SO4),  +  4H,O  +  2ND 

The  three  atoms  of  oxygen  furnished  ly  tiro  molecules  of  nitrate  suffice  to 
oxidize  three  atoms  of  copper;  so  that  3CuO  with  3HJ304 ,  may  form 
3CuS04  and  3H20  .  The  same  three  atoms  of  oxygen  (having  six  bonds) 
suffice  to  oxidize  six  molecules  of  ferrous  salt  into  three  molecules  of 
ferric  salt;  so  that  6FeS04  with  3H2S04 ,  can  form  3Fe2(S04)3  and  3H20  . 

Now  if,  by  the  last-named  reaction,  the  nitric  oxide  is  disengaged  in 
cold  solution,  with  excess  of  ferrous  salt  and  of  sulphuric  acid,  instead 
of  passing  off,  the  nitric  oxide  combines  with  the  ferrous  salt,  forming  a 
Wack-lrown  liquid,  (Fe304).,NO  ,  decomposed  by  heat  and  otherwise  un- 
stable: 2KN03  +  4H2S04  +  10FeS04  ==  K,S04  +  3Fe2(S04)s  +  4H20  + 
2(FeS04)2NO  . 

a. — This  exceedingly  delicate  "Brown  ring"  test  for  nitric  acid  or 
nitrates  in  solution  may  be  conducted  as  follows:  If  the  solution  of  a 
nitrate  is  mixed  with  an  equal  volume  of  concentrated  H2S04 ,  the  mixture 


CALIFORNIA   COLlEfii 

§241,  sd.  NITRIC  ACID.        &  PHARMAC¥9 

allowed  to  cool  and  a  concentrated  solution  of  FeSO^  then  cautiously  added 
to  it,  so  that  the  fluids  do  not  mix,  the  junction  shows  at  first  a  purple, 
afterwards  a  brown  color  (Fresenius,  QnaL  Anal,  IQtli  ed.,  387).  A  second 
method  of  obtaining  the  same  brown  ring  is:  Take  sulphuric  acid  to  a 
quarter  of  an  inch  in  depth  in  the  test-tube ;  add  without  shaking  a  nearly 
equal  bulk  of  a  solution  of  ferrous  sulphate,  cool;  then  add  slowly  of  the 
solution  to  be  tested  for  nitric  acid,  slightly  tapping  the  test-tube  on  the 
side  but  not  shaking  it.  The  brown  ring  forms  between  the  two  layers  of 
the  liquid.  A  third  method  often  preferred  is:  Take  ferrous  sulphate 
solution  to  half  ;in  inch  in  depth  in  the  test-tube;  add  two  or  three  drops 
of  the  liquid  under  examination  and  mix  thoroughly;  incline  the  test-tube 
and  add  an  equal  volume  of  concentrated  H2S04  in  such  a  way  that  it  will 
pass  to  the  bottom  and  form  a  separate  layer.  Cool  and  let  it  stand  a 
few  minutes  without  shaking.  Nitrous  acid  interferes  with  this  test  but 
the  brown  ring  is  produced  when  a  nitrite  is  acidified  with  acetic  acid 
while  sulphuric  acid  is  required  in  the  case  of  a  nitrate  (§239,  8).  If 
the  presence  of  a  nitrite  is  suspected,  the  solution  should  first  be  acidi- 
fied with  acetic  acid  and  ferrous  sulphate  added.  If  a  brown  color  is 
produced  a  nitrite  is  present.  The  nitrite  may  be  removed  by  boiling 
the  solution  until  the  brown  color  has  disappeared  and  does  not  return 
on  adding  more  ferrous  sulphate  solution  and  acetic  acid.  After  cool- 
ing the  solution,  a  nitrate  may  be  tested  for  by  means  of  strong  sul- 
phuric acid  and  ferrous  sulphate. 

If  strong  oxidizing  agents  are  present,  excess  of  FeS(>4  should  be  added 
and  the  solution  warmed.  The  solution  should  be  cooled  before  applying 
the  test.  Metals  forming  insoluble  sulphates  should  bo  removed  by  adding 
Na  C03  and  warming.  Iodides  interfere  by  forming  a  brown  ring  of  free 
iodine. 

6. — Indigo  solution. — In  presence  of  HC1  heat  moderately  and  the  blue 
color  is  destroyed.  Interfering  substances,  HC103 ,  HI03 ,  HBr03,  Fe'", 
CrVI ,  Mnvn ,  and  all  that  convert  HC1  into  C12 . 

c- — Sodium  salicylate  is  added  to  the  solution,  H2S04  is  slowly  added, 
the  test-tube  being  inclined.  Avoid  shaking,  keep  cool  for  five  minutes. 
A  yellow  ring  indicates  HN03  .  To  increase  the  brilliancy  of  the  color, 
shake,  cool  and  add  to  HN4OH  . 

d. — Ammonium  test. — Treat  the  solution  with  KOH  and  Al  wire,  warm 
until  gas  is  evolved.  Pass  the  gas  into  water  containing  a  few  drops  of 
Nessler's  reagent.  A  yellowish-brown  precipitate  indicates  HN03  : 
3HN03  +  8A1  -f  8KOH  =  3NHS  +  8KA102  +  H20  .  Nothing  interferes 
with  this  test,  but  action  is  delayed  by  Clv  ,  Iv  and  many  other  oxidisers. 
Ammonia  and  ammonium  salts  must  be  removed  by  evaporation  with 
before  applying  the  test. 


290  NITRIC  ACID.  §241,  80, 

e. — Nitrite  test. — Reduce  the  nitrate  to  nitrite  by  warming  with  Al  and 
KOH  .     At    short  intervals  decant  a  portion  of  the  solutioh,  add  a  drop  of 
KI,    acidify  with  HC-H302  and  test  for  I  with  CS2.     This  test  should 
always  be  made  in  connection  with  (d).     Other  oxidisers  including  Clv, 
Brv,  Iv,  and  Asv  are  reduced  before  the  reduction  of  the  HN03  begins: 
3HN03  +  2AI  +  5KOH  =  3KN02  +  2KA1O2  +  4H2O 
2KN02  +  2KI  +  4HC2H102  =  I2  +  4KC2H302  +  2H2O  +  2ND 

Other  means  of  making  the  nascent  hydrogen  are  sometimes  preferred; 
e.  g.,  sodium  amalgam,  a  mixture  of  Zn  and  Fe  both  finely  divided  and 
used  with  excess  of  hot  KOH ,  or  finely  divided  Mg  in  presence  of  H3P04 . 

/.  Various  organic  compounds  give  characteristic  color  reactions  with 
nitric  acid.  An  excellent  reagent  of  this  kind  consists  of  a  solution  of 
dimethyl  aniline  (2  drops)  and  p-toluidine  (0.2  grams)  in  50%  sul- 
phuric acid  (10  c.c.).  A  blood  red  color  is  produced  when  this  reagent 
is  brought  into  contact  with  even  a  dilute  solution  of  a  nitrate.  The 
color  appears  as  a  ring  test  between  the  two  liquids.  Reducing  agents 
especially  ferrous  salts  interfere  with  the  test.  By  the  addition  of  a 
crystal  of  KC103  and  HC1  and  boiling,  the  interference  may  be  overcome. 
KC103  gives  a  brown  color  but  on  dilution  the  red  color  of  the  nitrate 
generally  appears.  KC103  may  be  removed  by  heating  with  HC1.  (Wood- 
ruff, J.  Am.  Soc.  19,  156.  Schmidt  and  Lump  (Ber.  43,  794)  also 
give  a  color  reagent  which  gives  a  red  color.  The  reagent  is  a  solu- 
tion of  0.1  gram  of  Di-9,10-monoxyphenanthrylamine  in  1000  c.c. 
cone.  H2S04 .  The  dry  salt  or  the  residue  obtained  by  evaporation  of 
the  unknown  solution  is  added  to  a  few  c.c.  of  the  blue  reagent  which 
turns  red  in  the  presence  of  a  nitrate.  Oxidizing  agents  do  not  interfere. 

Add  three  drops  of  the  solution  to  be  tested  to  two  drops  of 
diphenylamine,  (C0H5)2NH,  dissolved  in  H2S04 .  A  blue  color  indicates 
a  nitrate.  Cl°,  Clv,  Brv,  Iv,  Mnvn,  Crvl,  Se1^,  and  Fe'"  interfere  with  this 
test.  This  test  is  of  especial  value  in  showing  the  absence  of  nitrates. 
If  no  color  is  obtained,  it  is  certain  that  no  nitrate  is  present. 

g. — Brucine,  dissolved  in  concentrated  sulphuric  acid,  treated  (on  a  porcelain 
surface)  with  even  traces  of  nitrates,  gives  a  fine  deep-red  color,  soon  paling  to 
reddish-yellow.  If  now  stannous  chloride,  dilute  solution,  be  added,  a  fine  red- 
violet  color  appears.  (Chloric  acid  gives  the  same  reaction.) 

h. — Phenol,  C6H5OH  ,  gives  a  deep  red-brown  color  with  nitric  acid,  by  for- 
mation of  nitrophenol  (mono,  di  or  tri),  C6H4(NO2;OH  to  C6H2(NO2)3OH , 
"picric  acid"  or  nitrophenic  acid.  A  mixture  of  one  part  of  phenol  (cryst. 
carbolic  acid),  four  parts  of  strong  sulphuric  acid,  and  two  parts  of  water, 
constitutes  a  reagent  for  a  very  delicate  test  for  nitrates  (or  nitrites) ,  a  few  drops 
being  sufficient.  With  unmixed  nitrates  the  action  is  explosive,  unless  upon  very 
small  quantities.  Ths  addition  of  potassium  hydroxide  d  epens  and  brightens  the 
c  lor.  According  to  Sprengel  (J.  C.,  1863,  15,  396),  the  somewhat  similar  o  l--r 
given  by  compounds  of  chlorine,  bromine,  iodine  and  by  organic  m  ttter  may  be 
removed  by  adding  ammonium  hydroxide  without  diminishing  the  brightness  of 
th  color  formed  by  the  nitrates. 


§242,  3.  OXYGEN.  291 

i. — According  "to  Lindo  (C.  N.,  1888,  68,  176),  resorcinal  is  five  times  more 
delicate  a  lest  than  phenol.  T(  n  grammes  of  r<  s;,rcinol  are  dissolved  in  100  cc. 
of  water;  one  dr  <p  of  this  solution  with  one  drop  of  a  15  per  cent  solution  of 
HC1  and  two  drops  of  concentrated  H2SO4  are  added  to  0.5  cc.  of  the  nitrate 
to  be  tested.  Nitrous  aeid  gi\es  the  same  purple  color. 

/.  A  little  pyrogallol  is  dissolved  iu  the  liquid  to  be  tested  (less  Ihan  one 
mg.  to  one  cc.)  and  ten  drops  of  concentrated  H./SO4  are  dropped  down  tin- 
side  of  the  test  tube  so  as  to  form  two  layers;  at  the  surface  of  contact  a 
brown  or  yellow  coloration  appears  if  nitric  acid  is  present.  One  rag.  of 
nitric  acid  in  one  litre  of  potable  water  can  thus  be  detected  (Curtman,  Arrh. 
Pltnnn.,  issii.  223,  711). 

9.  Estimation. —  (a)  If  the  base  is  one  capable  of  readily  forming-  a  silicate, 
the  nitrate  is  fused  with  SiOo  and  estimated  by  the  difference  in  weight.  (/>)  By 
treating  with  hot  sulphuric  acid,  passing  the  distillate  into  BaC03  and  esti- 
mating the  nitric  acid  by  the  amount  of  barium  dissolved,  (c)  Treating  with 
Al  and  KOH  and  estimating  the  distillate  as  NH;!  .  (tf)  Neutralizing  the  free 
acid  with  ammonium  hydroxide,  and  after  evaporation  and  drying  at  115°, 
weighing  as  ammonium  nitrate,  (e)  In  presence  of  free  H2S04  a  ferrous  solu- 
tion of  known  strength  is  added  in  excess  to  the  nitrate  and  the  amount  of 
ferrous  salt  remaining  is  determined  by  a  standard  solution  of  potassium 
permanganate,  (f)  The  volume  of  hydrogen  generated  by  the  action  of  potas- 
sium hydroxide  upon  a  known  quantity  of  aluminum  is  measured;  and  the 
test  is  then  repeated  under  the  same  conditions,  but  in  presence  of  the  nitrate. 
The  difference  in  the  volume  of  the  hydrogen  obtained  represents  the  quantity 
of  NHn  that  has  been  formed. 

§242.   Oxygen.     0  =  16.000.     Usual  valence  two. 

1.  Properties. — A  colorless,   odorless  gas;    specific  gravity,   1.10535   (Rayleigh, 
Proc.    Roy.   Soc.,    1897,    204).     When   heated  it   diffuses   through   silver   tubing 
quite  rapidly  (Troost,  C.  r.,  1884,  98,  1427).     It  may  be  liquefied  by  cooling  the 
gas  under  great  pressure  and  then  suddenly  allowing  it  to  expand  under  reduced 
pressure.     It  boils  at    —  113°  under  50  atmospheres  pressure;    and  at     -  184° 
under   one   atmosphere   pressure    (Wroblewski,    C.   r.,    1884,    98,    304    and   982). 
Its  critical  temperature  is  about    —  118.8°,  and  the  critical  pressure  50.8  atmos- 
pheres.    Specific  gravity  of  the  liquid  at  -  181.4°,  1.124  (Olszewski,  M.,  1887,  8, 
73).     Oxvgen  is  sparingly  soluble  in  water  with  a  slight  increase    in    the   vol- 
ume   (Wmkler,    B.,    1889,    22,    1764);     slightly   soluble   in   alcohol    (Carius,    A., 
1855,  94,   134).     Molten  silver  absorbs  about  ten  volumes  of  oxygen,  giving  it 
up  upon  cooling   (blossoming  of  silver  beads)    (Levol,   C.  r.,   1852,   35,   63).     It 
transmits  sound  better  than  air  (Bender,  B.,  1873,  6,  665).     It  is  not  combustible, 
but  supports  combustion  much  better  than  air.     In  an  atmosphere  of  oxygen, 
a  glowing  splinter  bursts  into  a  flame;  phosphorus  burns  with  vivid  incandescence; 
also  an  iron  watch  spring  heated  with  burning  sulphur.     It  is  the  most  negative 
of  all  the  elements  except  fluorine;    it  combines  directly  or  indirectly  with  all 
the  elements  except  fluorine;  with  the  alkali  metals  rapidly  at  ordinary  temperature. 
The  combination  of  oxygen  with  -elements  or  compounds  is  termed  combustion 
or  oxidation.     The  temperature   at  which  the   combination  takes   place   varies 
greatly:    phosphorus  at  60°;    hydrogen  in  air  at  552°;    in  pure  oxygen  at  530° 
(Mallard  and  Le  Chatelier,   BL,   1883,    (2),   39,   2);    carbon  disulphide  at   149°; 
carbon  at  a  red  heat;  while  the  halogens  do  not  combine  by  heat  alone. 

2.  Occurrence. — The  rocks,  clay  and  sand  constituting  the  main  part  of  the 
earth's  crust  contain  from  44  to  48  per  cent  of  oxygen;    and  as  water  contains 
88.81  per  cent,   it  has  been   estimated   that  one-half   of   the   crust  is   oxygen. 
Except  in  atmospheric  air,  which   contains   about  2;;  per  cent  of  uncombined 
oxygen,  it  is  always  found  combined. 

3.  Formation. —  ((/)  By  igniting  HgO  .     (b)  By  heating  KC103  to  350°,  KC104 
is  produced  and  oxygen  is  evolved;  at  a  higher  temperature  the  KC1O4  becomes 
KC1  .     In  the  presence  of  MnO2   the  KC103  is  completely  changed  to  KC1  at 
200°,  without  forming  KC104  ,  the  Mn02  not  being  changed.     Spongy  platinum, 
CuO  ,  Fe,O3  ,  PbO2  ,  etc..  may  be  substituted  for  Mn02  (Mills  and  Donald,  J.  C., 
1882,  41, "l8;  Baudrimont,  AIII.  N.,  1S72,  103,  370).     Spongy  platinum,  ruthenium, 


OXYGEN.  §242,  3«. 

rhodium  and  indium  with  chlorine  water  or  with  hydrogen  peroxide  evolve 
oxygen.  The  spongy  ruthenium  acts  most  energetically  (Schoenbein,  A.  C/,., 
1866,  (4),  7,  103).  (c)  Action  of  heat  on  similar  salts  furnishes  oxygen;  e.g., 
KC10  and  KC1O2  form  KC1 ,  KBr03  forms  KBr  ,  KI03  and  KEO4  "form  KI  ' 
and  KN03  forms  KN02  (at  a  white  heat  K20  ,  NO  and  0  are  formed),  (d)  By 
the  action  of  heat  on  metallic  oxides  as  shown  in  the  equations  below,  (e)  By 
heating  higher  oxides  or  their  salts  with  sulphuric  acid.  Crvi  is  changed  to 
Cr'"  ,  Co"'  to  Co"  ,  Ni'"  to  Ni"  ,  Biv  to  Bi'"  ,  Fevi  to  Fe"'  ,  Pbiv  to  Pb"  ,  and 
Mn"+n  to  Mn";  in  each  case  a  sulphate  is  formed  and  oxygen  given  off: 

a.   2HgO  (at  500°)  =  2Hg  +  O2 

&.    lOKClOg  (at  350°)  =  6KC1O4  +  4KC1  +  3O2  (Teed,  J.  C.,  1887,  51,  283) 
2KC103  (at  red  heat)  =  2KC1  +  :.O3 
2KC103  +  nMn02   (at  200°)  =  nMnO2  +  2KC1  +  302 

c.  KC102  =  KC1  +  O2 
2KBr03  =  2KBr  -f  3O2 
2KI08  =  2KI  +  302 
KIO4  =  KI  +  202 
2KN03  =  2KN02  +  O2 

4KN02  (white  heat)  =  2K2O  +  4NO  +  O2 

d.  2Pb3O4  (white  heat)  =  GPbO  +  02 
2Sb2O5  (red  heat)  =  2Sb,04  +  02 
Bi205  (red  heat)  =  Bi2O,  +  02 
4Cr03  (about  200°)  =  2Cr,O3  +  302 

4K2Cr207  (red  heat)  —  2Cr,,O3  +  4K2Cr04  +  3O2 

6Fe203  (white  heat)  =  4Fe3O4  +  O2 

3Mn02  (white  heat)  =  Mn3O4  +  02 

6Co2O3  (dull-red  heat)  =  4Co3O4  -f  02 

2Ni203  (dull-red  heat)  =  4NiO  -f  O2 

2Ag20  (300°)  =  4Ag  -f  02 

2BaO2  (800°)  =  2BaO  +  O2 

e.  2K2Cr207  +  8H2SO4  =  4KCr(S04)2  +  302  +  8H2O 
4KMnO4  +  6H2S04  =  2K2S04  +  4MnSO4  +  502  +  6H2O 
2Pb3O4  -f  6H2S04  =  6PbSO4  +  6H20  +  02 

4.  Preparation. — (a)  By  heating  KC103  to  200°  in  closed  retorts  in  the  pres- 
ence  of   MnO2    or   Fe203  .     If   KC1O3    be   heated   alone,    higher   heat    (350°)    is 
required,  and  the  gas  is  given  off  with  explosive  violence.     About  equal  parts 
of  the  metallic  oxide  and  KC1O3  should  be  "taken.     (6)   BaO  heated  in  the  air 
to  550°  becomes  BaO2  ,  and  at  800°   is  decomposed  into  BaO  and   O  ,  making 
theoretically  a  cheap  process,     (c)  By  heating  calcium  plumbate.     The  calcium 
plumbate  is  regenerated  by  heating  in  the  air  (Kassner,  J.  (7.,  1894,  66,  ii,  89). 
(d)  By  passing  sulphuric  acid  over  red-hot  bricks:  2H2S04  =  2S02  +  2H,0  -f  O2; 
the  S02   is  separated  by  water,  and  after  conversion  into  H2SO4    (§266,  4)   is 
used  over  again,     (e)  By  warming  a  saturated  solution  of  chloride  of  lime  with 
a  small  amount  of  cobaltic  oxide,  freshly  prepared  and   moist.     The  cobaltic 
oxide  seems  to  plajr  the  same  role  as  NO  in  making  H2S04   (Fleitmaim,  A.  Ch., 
1865,  (4),  5,  507).    \f)  The  following  cheap  process  is  now  employed  on  a  large 
scale.     Steam  is  passed  over  sodium  manganate  at  a  dull-red  heat;  Mn2O,  and 
oxygen  are  formed.     Then,  without  change  of  apparatus  or  temperature,  air 
instead  of  steam  is  passed  over  the  mixture  of  Mn2O3  and  NaOH  .     The  Mn20, 
is  thus  again  oxidized  to  Na2Mn04  ,  and  free  nitrogen  is  liberated: 

4Na2MnO4  +  4H20  (dull-red  heat)  =  SNaOH  +  2Mn203  -f  30, 
8NaOH  +  2Mn203  +  air,  3(02  +  4N2)  =  4Na2MnO4  +  4H20  +  12N2 

5.  Solubilities.— See  1. 

6.  Reactions. — Pure  oxygen  may  be  breathed  for  a  short  time  without  injury. 
A  rabbit  placed  in  pure  oxygen  at  24°  liveol  for  three  weeks,  eating  voraciously 


§243. 


OZONE. 


all  the  tim^,  but  nevertheless  becoming  thin.  The  action  of  oxygen  at  7  2"  is 
to  produce  narcotism  and  eventually  death.  When  oxygen  is  cooled  by  a 
freezing"  mixture  it  induces  so  intense  a  narcotism  that  operations  may  be 
performed  under  its  influence.  Compressed  oxygen  is  "  the  most  fearful  poison 
known."  The  pure  gas  at  a  pressure  of  3.5  atmospheres,  or  air  at  a  pressure 
of  22  atmospheres,  produces  violent  convulsions,  simulating-  those  of  strychnia 
poisoning1,  ultimately  causing1  death.  The  arterial  blood  in  these  cases  is  found. 
to  contain  about  twice  the  quantity  of  its  normal  oxygen.  Further,  compressed 
oxygen  stops  fermentation,  and  permanently  destroys  the  power  of  yeast. 

At  varying'  temperatures  oxygen  combines  directly  with  all  metals  except 
silver,  gold  and  platinum,  and  with  these  it  may  be  made  to  combine  by  pre- 
cipitation. It  combines  with  all  non-metals  except  fluorine;  the  combination 
occurring-  directly,  at  high  temperatures,  except  with  Cl  ,  Br  and  I,  which 
require  the  intervention  of  a  third  body. 

7.  Ignition. — Most    elements    when    ignited    with    oxygen    combine    readily. 
Some  lower  oxides  combine  with  oxygen  to  form  higher  oxides,  and  certain 
other  oxides  evolve  oxygen,  forming  elements  or  lower  oxides.     Oxides  of  gold, 
platinum  and  silver  cannot  be  formed  by  igniting  the  metals  in  oxygen;  they 
must  be  formed  by  precipitation. 

8.  Detection. — Uncombined  oxygen  is  detected  by  its  absorption  by  an  alka- 
line  solution   of   pyrogallol;    by   the   combination    \vith   indigo   white   to   form 
indigo  blue;  by  its  combination  with  colorless  NO  to  form  the  brown  NO2;  by 
its  combination   with  phosphorus,   etc.     It  is  separated   from   other   gases  by 
its  absorption  by  a  solution  of  chromous  chloride,  pyrogallol  or  by  phosphorus. 
In  combination  in  certain  compounds  it  is  liberated  in  whole  or  in  part  by 
simple  ignition;  as  with  KC1O3  ,  KMn04  ,  HgO  ,  Au2O3  ,  Pt02  ,  Ag.O  ,  Sb,05  , 
etc.     In  other  combinations  by  ignition  with  hydrogen,  forming  water. 

9.  Estimation. — Free  oxygen  is  usually  estimated  by  bringing  the  gases^  in 
contact  with  phosphorus  or  with  an  alkaline  solution  of  pyrogallol  (CO,  having 
been  previously  removed),  and  noting  the  dimunition  in  volume.     Oxygen  in 
combination  is  usually  estimated  by  difference. 


§243.  Ozone.     0:i  =  48.000. 
0  —  0 


Ozone  was  first  noticed  by  Van  Marum  in  1785  as  a  peculiar  smelling  gas 
formed  during  the  electric  discharge;  and  which  destroyed  the  lustre  of 
mercury.  Schoenbein  (Poyg.,  1840,  50,  616)  named  the  gas  ozone  and  noticed 
its  powerful  oxidizing  properties.  It  is  said  to  be  an  ever-present  Constituent 
of  the  air,  giving  to  the  sky  its  blue  color;  present  much  more  in  the  country 
and  near  the  seashore  than  in  the  air  of  cities  (Hartley,  J.  C.,  1881,  39,  57  and 
111;  Houzeau,  C.  r.,  1872,  74,  712).  Ozone  is  always  mixed  with  ordinary  oxygen, 
partly  due  to  dissociation  of  the  ozone  molecule,  which  is  stable  only  at  low 
temperatures  (Hautefeuille  and  Chappuis,  C.  r.,  1880,  91,  522  and  815).  It  is 
prepared  by  the  action  of  the  electric  discharge  upon  oxygen  (Bichat  and 
Guntz,  C.  r.,  1888,  107,  344;  Wills,  /?.,  1873,  6,  769).  By  the  oxidation  of  moist 
phosphorus  at  ordinary  temperature  (Leeds,  A.,  1879,  198,  30;  Marig-nac,  C.  r., 
1845,  20,  808).  By  electrolysis  of  dilute  sulphuric  acid,  using  lead  electrodes 
Planti,  C.  r.,  1866,  63,  181).  By  the  action  of  concentrated  sulphuric  acid  on 
potassium  permanganate  (Schoenbein,  «/.  pr.,  1862,  86,  70  and  377).  Many 
readily  oxidized  organic  substances  form  some  ozone  in  the  process  of  oxida- 
tion (Belluci,  B.,  1879,  12,  1699).  Ozone  is  a  gas,  the  blue  color  of  which  can 
be  plainly  noticed  in  tubes  one  metre  long.  Its  odor  reminds  one  somewhat 
of  chlorine  and  nitrogen  peroxide,  noticeable  in  one  part  in  1,000,000.  In  strong 
concentrations  it  acts  upon  the  respiratory  organs,  making  breathing  difficult. 
When  somewhat  concentrated  it  attacks  the  mucous  membrane.  It  caused 
death  to  small  animals  which  have  been  made  to  breathe  it.  For  further  con- 
cerning the  physiological  action,  see  Binz,  C.  C.,  1873,  72.  Its  specific  gravity 


294  HYDROGEN   PEROXIDE.  §244. 

is  1.658  (Soret,  A.,  1866,  138,  4).  It  has  been  liquefied  to  a  deep-blue  liquid, 
boiling  at  -  106°  (Olszewski,  M.,  1887,  8,  230).  The  gas  is  sparingly  soluble 
in  water  (Carius,  B.,  1873,  6,  806).  It  decomposes  somewhat  into  inactive  oxygen 
at  ordinary  temperature,  and  completely  when  heated  above  300°,  with  increase 
of  volume.  A  number  of  substances  decompose  ozone,  without  themselves 
being  changed;  e.  g.,  platinum  black,  platinum  sponge,  oxides  of  gold,  silver, 
iron  and  copper,  peroxides  of  lead,  and  manganese,  potassium  hydroxide,  etc. 
It  is  one  of  the  most  active  oxidizing  agents  known,  the  presence  of  water  being 
necessary.  When  ozone  acts  as  an  oxidizing  agent,  there  is  no  change  in  volume; 
but  one-third  of  the  oxygen  entering  into  the  reaction,  inactive  oxygen  remaining. 

Moist  ozone  oxidizes  all  metals  except  gold  and  platinum  to  the  highest  pos- 
sible oxides. 

Pb"  becomes  PbO2 

Sn"  becomes  Sn02 

Hg1'  becomes  Hg" 

Bi'"  becomes  BL05 

Pd"  becomes  Pd02 

Cr'"  becomes  Crvi 

Fe"  becomes  Fe,O3;  in  presence  of  KOH  ,  K,Fe04 

Mn"  becomes  MnO,;  in  presence  of  H2SO4  or  HNO3  ,  HMnO4  is  formed. 

Co"  becomes  Co"' 

Ni"  becomes  Ni'"  .  With  the  salts  of  nickel  and  cobalt  the  action  is  slow, 
rapid  with  the  moist  hydroxides. 

K4Fe(CN)6  becomes  K3Fe(CN)6 

N203  becomes  HN03  ,  in  absence  of  water  N02  is  formed 

SO2  becomes  H2SO4 

H,S  becomes  S  and  H,O  ,  the  sulphur  is  then  oxidized  to  H2SO4  (Pollacci, 
C.  C.,  1884,  484) 

P  and  PH3  become  H3P04 

HC1  becomes  Cl  and  H,O 

HBr  becomes  Br  and  ELO 

I  becomes  HIO3  and  HI04  (Ogier,  (7.  r.,  1878,  86,  722) 

HI  and  KI  become  I  and  H20  ,  then  IV 

Most  organic  substances  are  decomposed;  indigo  is  bleached  much  more 
rapidly  than  by  chlorine  (Houzeau,  C.  r.,  1872,  75,  349). 

Alcohol  and  ether  are  rapidly  oxidized  to  aldehyde  and  acetic  acid. 

Ozone  is  usually  detected  by  the  liberation  of  iodine  from  potassium  iodide, 
potassium  iodide  starch  paper  being  used.  Because  HN02  and  many  other 
substances  give  the  same  reaction,  thallium  hydroxide  paper  is  preferred  by 
Si-hoene  (/?.,  1880,  13,  1508).  The  paper  is  colored  brown,  but  the  reaction  is 
much  less  delicate  than  with  potassium  iodide  starch  paper.  Chlorine,  bromine, 
iodine  and  nitrous  oxides  do  not  interfere  with  the  following  test.  Paper  is  moistened 
with  a.  15  per  cent  solution  of  potassium  iodide  to  which  a  1  per  cent  alcoholic 
rosolic  acid  or  phenolphthalein  solution  has  been  added  until  a  marked  opalescenco 
is  produced.  A  red  color  is  produced  by  exposure  to  ozone.  The  color  produced 
by  the  rosolic  acid  is  more  permanent  than  the  phenolphthalein  color.  It  i  ? 
estimated  quantitatively  by  passing  the  gas  through  a  neutral  solution  of  KI  and 
titration  of  the  liberated  iodine:  O3  +  2KI  +  H2O  =  O2  +  I2  +  2KOH  . 

Ozone  may  also  be  detected  by  the  blue  color  imparted  to  guaiacum  tincture. 
Hydrogen  peroxide  does  not  produce  this  effect  unless  ferrous  sulphate  solution  is 
added.  The  most  sensitive  reagent  is  a  freshly  prepared  10  per  cent  solution  of 
guaiacum  gum  in  50  per  cent  water  solution  of  chloralhydrate.  (Weber,  Z.,  43,  47). 

§244.  Hydrogen  peroxide.     H202  =  34.016. 
H  — 0  — 0— H. 

1.  Properties. — Pure  hydrogen  peroxide  (99.1  per  cent)  is  a  colorless  syrupy 
liquid,  boiling  at  84°  to  85°  at  68  mm.  pressure.  It  does  not  readily  moisten 
the  containing  vessel.  It  is  volatile  in  the  air,  irritating  to  the  skin,  and 


$244,  6,  10.  HYDROGEN   PEROXIDE.  295 

reacts  strongly  acid  to  litmus.  The  ordinary  three  per  cent  solution  ean  he 
evaporated  on  the  water  bath  until  it  contains  about  HO  per  eent  H2O2  ,  losing 
about  one-half  by  volatilization.  The  presence  of  impurities  causes  its  decom- 
position with  explosive  violence.  Before  final  concentration  under  reduced 
pressure  it  should  be  extracted  with  ether  (Wolffenstein,  J?.,  1894,  27,  iJiiOT). 
The  dilute  solutions  are  valuable  in  surgery  in  oxidizing  putrid  flesh  of  wounds, 
etc.;  they  are  quite  stable  and  may  be  preserved  a  long  time  especially  if  acid 
(Hanriott,  C.  r.,  1885,  100,  57).  The  presence  of  alkalis  decreases  the  stability. 
Concentrated  solutions  evolve  oxygen  at  20°,  and  frequently  explode  when 
heated  to  nearly  100°.  It  contains  the  most  oxygen  of  any  known  compound; 
one-half  of  the  oxygen  being  available,  the  other  half  combining  with  the 
hydrogen  to  form  water. 

2.  Occurrence. — In  rain  water  and   in  snow    (Houzeau,   C.  r.,   1870,  70,   519). 
It  is  also  said  to  occur  in  the  juices  of  certain  plants. 

3.  Formation. —  (or)  By  the  electrolysis  of  70  per  cent  H,SOt   (Richarz,  W.  A., 
1887,  31,  912).     O)   By  the  action  of  ozone  upon  ether  and  water    (Berthelot, 
C.  r.,  1878,  86,  71).     (c)  By  the  action  of  ozone  upon  dilute  ammonium  hydroxide 
(Carius,  B.,  1874,  7,  1481).     (d)  By  the  decomposition  of  various  peroxides  with 
acids,     (e)  By  the  action  of  oxygen  and  water  on  palladium  sponge  saturated 
with  hydrogen  (Traube,  B.,  1883,  16,  1201).     (f)  By  the  action  of  moist  air  on 
phosphorus  partly  immersed  in  water  (Kingzett,  J.  C.,  1880,  38,  3). 

4.  Preparation.— BaO2    is    decomposed    by    dilute    H2SO,  ,    the    BaSOt    being 
removed  by  filtration.     The  BaO,  is  obtained  by  heating  BaO  in  air  or  oxygen 
to  low  redness.     At  a  higher  heat  the  Ba.O.,   is  decomposed   into  BaO   and   O 
(Thomsen,   B.,   1874,   7,   73).     Sodium  peroxide,   Na,0,  ,   is   formed   by   heating 
sodium  in  air  or  oxygen   (Harcourt,  J.  C.,   1862,   14,  267);  by  adding  H2O2   to 
NaOH  solution  and  precipitating  with  alcohol.     Prepared  by  the  latter  method 
it  contains  water. 

5.  Solubilities It    is   soluble   in    water   in   all   proportions:   also   in    alcohol, 

which  solvent   it  slowly  attacks.     Ba02   is  insoluble  -in  water,  decomposed    by 
acids,  including  CO,  and  H,SiFR  with  formation  of  H2O,  .     Na.2O2  is  soluble  in 
water  with  generation  of  much  heat.     It  is  a  powerful  oxidizing  agent. 

0.  Reactions.     A. — With    metals     and    their    compounds. — Hydrogen 
peroxide  usually  acts  as  a  powerful  oxidizing  agent  to  the  extent  of  one- 
half  its  oxygen.     Under  certain  conditions,  however,  it  acts  as  a  strong 
reducing  agent.     Some  substances  decompose  it  into  H20  and  0  without 
changing  the  substance  employed,  e.  g.,  gold,  silver,  platinum,  manganese 
dioxide,   charcoal,  etc.   (Kwasnik,  B.,   1892,   25,   67).     Many  metals  are 
oxidized  to  the  highest  oxides,  e.  g.,  Al ,  Fe  ,  Mg  ,  Tl ,  As  ,  etc.     Gold  and 
platinum  are  not  attacked. 

1.  Pb"  becomes  Pb02  (Schoenbein,  J.  pr.,  1862,  86,  129;  Jannasch  and 
Lesinsky,  B.,  1893,  26/2334). 

2.  Ag20  becomes  Ag  and  0  . 

3.  HgO  becomes  Hg  and  0  . 

4.  Au.,0.,  becomes  An  and  0  . 

5.  As"'  becomes  Asv. 

6.  Sn"  becomes  Sniv. 

7.  Bi"'  becomes  Biv. 

8.  Cu"  in  alkaline  solution  (Fehling's  solution)  becomes  Cu20  (Hanriott. 
BL,  1886,  (2),  46,  468). 

9.  Fe"  becomes  Fe"'  (Traube,  B.,  1884,  17,  1062). 

10.  Tl'  becomes  T1203  (Schoene,  A.,  1879,  196,  98). 


296  HYDROGEN   PEROXIDE.  §244,  6,  11. 

11.  Cr'"  becomes  CrVI  in  alkaline  mixture  (Lenssen,  J.pr.y  1860,  81,  278). 

12.  Crvl  with  H2S04  gives  a  blue  color,  HCr04 ,  perchromic  acid,  soon 
changing  to  green  by  reduction  to  Cr'".     By  passing  the  air  or  vapor 
through  a  chromic  acid  solution,  ozone  is  separated  from  hydrogen  perox- 
ide, the  latter  being  decomposed  (Engler  and  Wild,  B.9  1896,  29,  1940). 

13.  Mn"  in  alkaline  mixture  becomes  Mn02 .     In  presence  of  KCN  a 
separation  from  Zn  (Jannasch   and  Niederhofheim,  B.,  1891,  24,   3915; 
Jannasch,  Z.  anorg.,  1896,  12,  124  and  134). 

Mn"+n  with  H2S04  forms  MnS04 ,  oxygen  being  evolved  both  from  the 
H202  and  from  the  Mn  compound  (Brodie,  J.  C.,  1855,  7,  304;  Lunge, 
Z.  angew.,  1890,  6). 

14.  BaO ,  SrO ,  and  CaO  become  the  peroxides. 

15.  NaOH  becomes  Na202.8H20  . 

73.  NH4OH  becomes  NH4N02  (Weith  and  Webber,  B.,  1874, 7,  17  and  45). 
17.  TiIV  is  oxidized  to  pertitanic  acid,  H2Ti04,  with  the  production  of  ;i 
yellow  color  which  constitutes  a  very  delicate  test  for  H202 . 
B. — With  non-metals  and  their  compounds. 

1.  K4Fe(CN)6  becomes  K3Fe(CN)f,   (Weltzien,  A.,   1866,   138,   129);  in 
alkaline  solution  the  reverse  action  takes  place:    2K3Fe(CN)6  -f-  2KOH  + 
H202  =  2K4Fe(CN)6  +  2H20  +  02  (Baumann,  Z.  angew.,  1892,  113). 

2.  03  becomes  02  (Schoene,  I.  c.,  page  239). 

3.  H3P02  becomes  H3P04 . 

4.  H2S  and  sulphides,  and  S02  and  sulphites,  become  H2S04  or  sulphates 
(Classen  and  Bauer,  B.,  1883,  16,  1061). 

5.  Cl  becomes  HC1  (Schoene,  /.  c.,  page  254).     It  is  a  valuable  reagent 
for  the  estimation  of  chloride  of  lime :  CaOCl2  +  H202  =  CaCl2  +  H20  + 
02  (Lunge,  Z.  angew.,  1890,  6). 

6.  I  becomes  HI  (Baumann,  Z.  angew.,  1891,  203  and  328).     KC1 ,  KBr  , 
and  KI  liberate  oxygen  from  H202  but  no  halogen  is  set  free;  except  that 
with   commercial   H202   free   iodine   may   always   be   obtained   from   KI 
(Schoene,  A.,  1879,  195,  228;  Kingzett,  J.  C.,  1880,  37,  805). 

7.  Ingition. — The  peroxide  of  barium  is  formed  by  igniting-  BaO  to  dull  red- 
ness; strong-  ig-nition  causes  decomposition  of  the  BaO2  into  BaO  and  O  .     The 
peroxide  of  calcium  cannot  be  formed  by  ignition  of  lime  in  air  or  oxygen. 

8.  Detection. — In  a  dilute   solution   of  tincture   of   guaiac   mixed  with  malt 
infusion  or  ferrous  sulphate,  a  blue  color  is  obtained  when  H2O2  is  added.     To 
the  solution  supposed  to  contain  I^Oa  add  a  few  drops  of  lead  acetate;    then 
KI ,  starch,  and  a  little  acetic  acid;    with  H2O2  a  blue  color  is  produced  (Schoen- 
bein,  I.  c.;    Struve,  Z.,   1869,  8,  274).     As  confirmatory,  its  action  on  KMnO4 
and  on   K2Cr2O7  should  be  observed.     A  ten  per  cent  solution  of  ammonium 
molybdate  with  equal  parts  of  concentrated  sulphuric  acid  gives  a  characteristic 
deep  yellow  color  with  H2O2  (Deniges,  C.  r.,  1890,  110,  1007;    Crismer,  BL,  1891, 
(3),   6,    22).     H2O2   gives   some   extremely   delicate   color   tests   with   the   analine 
bases  (Ilosyay,  B.,  1895,  28,  2029;   Deniges,  /.  Pharm.,  1892,  (5),  25,  591 ;  Weber,  Z., 
43,  47).    Titanium  sulphate,  Ti(SO4)2,  in  acid  solution  gives  an  extremely  delicate 
test  for  H2O2 ,  a  yellow  color  being  produced.     On  the  addition  of  caustic  alkali  a 
yellowish  orange  precipitate  is  produced  which  redissolves  in  excess  of  the  reagent. 


£245,  2.  FLUORINE.  297 

9.  Estimation. — (a)  By  measuring  the  amount  of  oxygen  liberated  with  MnO 
(Hanriott,  BL,  1885,  (2),  43,  468).  (ft)  "By  the  amount  of  standard  KMnO4  reduced, 
or  by  measuring  the  volume  of  oxygen  set  free:  2KMnO4  -f  3H2SO4  +  5H2O2  = 
K2SO4  +  2MnSO4  +  8H2O  -f-  5O2 .  (c)  By  decomposition  of  KI  in  presence  of  an 
excess  of  dilute  H2SO4 ;  and  titration  of  the  liberated  iodine  with  standard  Na2S2O3 . 
(t/)  Dissolve  a  weighed  sample  of  BaO2  in  dilute  HC1 ,  add  K3Fe(CN)6 ;  transfer  to 
an  azotometer  and  add  KOH .  The  volume  of  oxygen  is  a  measure  of  the  amount 
of  H2O2  (Baumann,  I.  c.). 


§245.  Fluorine.     F  =  19.00.     Valence  one. 

Since  Davy's  experiments  in  1813,  many  others  have  attempted  the  isolation 
of  fluorine.  In  his  zeal  the  unfortunate  Lonyet  fell  a  victim  to  the  poisonous 
fumes  which  he  inhaled.  Faraday,  (lore,  Fremy,  and  others  took  up  the  prob- 
lem in  succession,  but  it  was  not  ultimately  solved  until  H.  Moissan,  in  1886, 
produced  a  gas  which  the  chemical  section  of  the  French  Academy  of  Sciences 
decided  to  be  fluorine.  Many  ingenious  experiments  had  been  made  in  order 
to  obtain  fluorine  in  a  separate  state,  but  it  was  found  that  it  invariably 
combined  with  some  portion  of  the  material  of  the  vessel  in  which  the  opera- 
tion was  conducted.  The  most  successful  of  the  early  attempts  to  isolate 
fluorine  appears  to  have  been  made,  at  the  suggestion  of  Davy,  in  a  vessel  of 
fluor-spar  itself,  which  could  not,  of  course,  be  supposed  to  be  in  any  way 
affected  by  it.  Moissan's  method  was  as  follows:  The  hydrofluoric  acid  having 
been  very  carefully  obtained  pure,  a  little  potassium  hydrofluoride  was  dis- 
solved in  it  to  improve  its  conducting  power,  and  it  was  subjected  to  the  action 
of  the  electric  current  in  a  U  tube  of  platinum,  down  the  limbs  of  which  the 
electrodes  were  inserted;  the  negative  electrode  was  a  rod  of  platinum,  and 
the  positive  was  made  of  an  alloy  of  platinum  with  10  per  cent  of  iridium.  The 
U  tube  was  provided  with  stoppers  of  fluor-spar,  anct "platinum  delivery  tubes 
for  the  gases,  and  was  cooled  to  — 23°.  The  gaseous  fluorine,  which  was  extri- 
cated at  the  positive  electrode,  was  colorless,  and  possessed  the  properties  of 
chlorine,  but  much  more  strongly  marked.  It  decomposed  water  immediately, 
seizing  upon  its  hydrogen,  and  liberating  oxygen  in  the  ozonized  condition;  it 
exploded  with  hydrogen,  even  in  the  dark,  and  combined,  with  combustion, 
with  most  metals  and  non-metals,  even  with  boron  and  silicon  in  their  crystal- 
lized modifications.  As  ,  Sb  ,  S  ,  I ,  alcohol,  ether,  benzol  and  petroleum  took 
fire  in  the  gas.  Carbon  was  not  attacked  by  it  (Moissan,  1886,  C.  r.,  103,  202 
and  256;  J.  C.,  50,  1886,  849  and  976;  A.  Ch.,  1891,  (6),  24,  224). 

1.  Properties. — A  gas  of  light  greenish-yellow  color  and  strong  pungent  odor; 
Specific  gravity,    1.313    (Moissan,    C.   R.,    109).     When  cooled  to   a   temperature 
of  about     -  187°  it  condenses  to  a  mobile  yellow  liquid.     Specific   gravity  of 
this  liquid  is  1.14.     At  —  223°  C.,  fluorine  solidifies  to  a  pale  yellow  solid.     The 
solid  loses  its  yellow  color  and  becomes  perfectly  white  at    —  252°    (Moissan 
and  Dewar,  C.  R.,  1903,  136). 

2.  Occurrence. — Fluorine  does  not  occur  free  in  nature,  but  is  found  in  con- 
siderable quantities  in  combination  with  calcium  in  the  mineral  fluorspar,   CaF2 . 
Its  other  fairly  common  compounds  are  cryolite,  Na3AlF6  and  apatite,  Ca5(PO4)sF  . 

Fluorine,  in  several  characteristics,  appears  as  the  first  member  of  the 
Chlorine  Series  of  Elements.  It  cannot  be  preserved  in  the  elemental  state, 
as  it  combines  with  the  materials  of  vessels  (except  fluor-spar),  and  instantly 
decomposes  water,  forming  hydrofluoric  acid,  HF  ,  an  acid  prepared  by  acting 
on  calcium  fluoride  with  sulphuric  acid  (a).  Fluorine  also  combines  with 
silicon,  forming  SiF4  ,  silicon  fluoride,  a  gaseous  compound,  generally  prepared 
by  the_  action  of  concentrated  sulphuric  acid  on  calcium  fluoride  and  silicic 
anhydride.  (ft).  On  passing  silicon  fluoride  into  water,  a  part  of  it  is  transposed 
by  the  water,  forming  silicic  and  hydrofluoric  acids  (c);  but  this  hydrofluoric 
acid  does  not  remain  free,  but  combines  with  the  other  part  of  the  fluoride  of 
silicon,  as  fluosilicic  acid  (hydrofluosilicic  acid],  (HF)2SiF4  or  H2SiF6  (d)  (Offer- 
mann,  Z.  angeiv.,  1890,  617).  This  acid  is  used  as  a  reagent;  forming  metallic 
fluosilicates  (silicofluorides),  soluble  and  insoluble  (§246). 


298  HYDROFLUORIC    ACID—FLUOSILICIC   ACID.  §246. 

a.  CaF2  +  H2SO4  =  CaSO4  +  2HF 

6.  2CaF2  +  SiO2  +  2H2SO4  =  2CaSO4  +  2H2O  +  SiF4 

c.  SiF4  +  4H2O  =  Si(OH)4  +  4HF  (not  remaining  free) 

d.  2HF  +  SiF4  »  H2SiF6 

c  and  d.     3SiF4  +  4H2O  =  Si(OH)4  +  2H2SiF6 

§246.  Hydrofluoric  acid.     HF  =  20.008. 
H'F-',  H  —  F  . 

A  colorless,  intensely  corrosive  gas,  soluble  in  water  to  a  liquid  that  reddens 
litmus,  rapidly  corrodes  glass,  porcelain,  and  the  metals,  except  platinum  and 
gold  (lead  but  slightly).  Both  the  solution  and  its  vapor  act  on  the  flesh  as 
an  insidious  and  virulent  caustic,  giving  little  warning,  and  causing  obstinate 
ulcers.  The  anhydrous  acid  at  25°  has  a  vapor  density  of  20,  indicating  that 
the  molecule  at  this  temperature  is  H2F2  .  But  at  100°  it  is  only  10,  indicating 
that  at  that  temperature  the  molecule  is  HF  .  The  anhydrous  liquid  acid 
boils  at  19.44°  and  does  not  solidify  at  — 34.5°. 

The  fluorides  of  the  alkali  metals  are  freely  soluble  in  water,  the  solutions 
alkaline  to  litmus  and  slightly  corrosive  to  glass;  the  fluorides  of  the  alkaline 
earth  metals  are  insoluble  in  water;  of  copper,  lead,  zinc  and  ferricum,  spar- 
ingly soluble;  of  silver  and  mercury  readily  soluble.  Fluorides  are  identified 
by  the  action  of  the  acid  upon  glass. 

Calcium  chloride  solution  forms,  in  solution  of  fluorides  or  of  hydrofluoric 
acid,  a  gelatinous  and  transparent  precipitate  of  calcium  fluoride,  CaF,  ,  slightly 
soluble  in  cold  hydrochloric  or  nitric  acid  and  in  ammonium  chloride  solution. 
Barium  chloride  precipitates,  from  free  hydrofluoric  acid  less  perfectly  than 
from  fluorides,  the  voluminous,  white,  barium  fluoride,  BaF.  .  Silver  nitrate 
gives  no  precipitate. 

Sulphuric  acid  transposes  fluorides,  forming  hydrofluoric  acid.  HF  (§245,  a). 
The  gas  is  distinguished  from  other  substances  by  etching  hard  glass— previously 
prepared  by  coating  imperviously  with  (melted)  wax,  and  writing  through  the 
coat.  The  operation  may  be  conducted  in  a  small  leaden  tray,  or  cup  formed 
of  sheet  lead;  the  pulverized  fluoride  being  mixed  with  sulphuric  acid  to  the 
consistence  of  paste. 

If  the  fluoride  be  mixed  with  silicic  acid,  we  have,  instead  of  hydrofluoric 
acid,  silicon  fluoride,  SiF4  (§245,  &);  a  gas  which  does  not  attack  glass,  but  when 
passed  into  water  produces  fluosilicic  acid,  H2SiF6  (§245,  c  and  d).  See  below. 

Also,  heated  with  acid  sulphate  of  potassium,  in  the  dry  way,  fluorides  dis- 
engage hydrofluoric  acid.  If  this  operation  be  performed  in  a  small  test-tube, 
the  surface  of  the  glass  above  the  material  is  corroded  and  roughened:  CaF.,  + 
:2KHS04  =  CaS04  +  K2SO4  +  2HF  .  By  heating  a  mixture  of  borax,  acid 
sulphate  of  potassium,  and  a  fluoride,  fused  to  a  bead  on  the  loop  of  platinum 
wire,  in  the  clear  flame  of  the  Bunsen  gas-lamp,  an  evanescent  yellowish-green 
color  is  imparted  to  the  flame. 

§247.  Fluosilicic  acid.     H2SiF6  =  144.316. 

Fluosilicic  acid  *  (hydrofluosilicic  acid),  (HF)2SiF4,  or  H2SiF6,  is  soluble  in  water 
and  forms  metallic  fluosilicates  (silico fluorides] ,  mostly  soluble  in  water;  those  of 
barium  (§186,  6i),  sodium  and  potassium,  being  only  slightly  soluble  in  water,  and 
made  quite  insoluble  by  addition  of  alcohol. 

*  Fluosilicic  acid  is  directed  to  be  prepared  by  taking  one  part  each  of  fine  sand  and  finely  pow- 
dered fluor-spar,  with  six  to  eight  parts  of  concentrated  sulphuric  acid,  in  a  small  stoneware 
bottle  or  a  glass  flask,  provided  with  a  wide  delivery-tube,  dipping  into  a  little  mercury  in  a 
small  porcelain  capsule,  which  is  set  in  a  large  beaker  containing  six  or  eight  parts  of  water. 
The  stoneware  bottle  or  flask  is  set  in  a  small  sand-bath,  with  the  sand  piled  about  it,  as  high  as 
the  material,  and  gently  heated  from  a  lamp.  Each  bubble  of  gas  decomposes  with  deposition 
of  gelatinous  silicic  acid.  When  the  water  is  filled  with  this  deposit,  it  may  be  separated  by 
straining  through  cloth  and  again  treating  with  the  gas  for  greater  concentration.  The  strained 
liquid  is  finally  filtered  and  preserved  for  use. 


§249,  4.  SILICON— SILICON  DIOXIDE.  299 

Potassium  fluosilicate  is  precipitated  translucent  and  gelatinous.  Ammonium 
hydroxide  precipitates  silicic  acid  with  formation  of  ammonium  fluoride.  With 
roncentra  ed  sulphuric  acid,  they  disengage  hydrofluoric  acid,  HF.  By  heat,  they 
are  resolved  into  fluorides  and  silicon  fluoride:  BaSiF6  =  BaF2  +  SiF4. 

§248.  Silicon.     Si  =  28.3.     Valence  four  (§15). 

There  are  three  modifications  of  silicon:  («)  Amorphous. — A  dark  brown 
powder;  specific  <irurity,  2.0;  non-volatile;  infusible;  burns  in  the  air,  forming 
SiO2  ,  and  in  chlorine,  forming-  SiCl4  .  It  is  not  attacked  by  acids  except  HF: 
Si  +  6HF  =  H2SiF0  +  2H2  .  It  is  dissolved  by  KOH  \vith  evolution  of 
hydrogen.  (&)  Grapliitoulal. — May  be  fused,  but  is, not  oxidized  upon  ignition 
in  air  or  in  oxygen.  It  is  not  attacked  by  HF  ,  but  is  dissolved  by  a  mixture 
of  HF  and  HNO3  ,  forming-  H2SiFo  .  It' is  attacked  slowly  by  fused  KOH  . 
(c)  AddiiHintitH'  silicon,  crystalline  silicon. — Grayish-black,  lustrous,  octahedral 
crystals,  formed  by  fusing-  the  graphitoidal  form.  Specific  gravity,  2.49  at  10° 
(Woehler,  A.,  1856,  97,  261).  It  scratches  glass  but  not  topaz.  It  melts  between 
the  melting-  points  of  pig-  iron  and  steel,  1100°  to  1300°.  In  chemical  properties 
it  is  very  similar  to  the  graphitoidal  form,  being  attacked  with  even  greater 
ditliculty.  Silicon  is  never  found  free  in  nature,  but  always  in  combination  as 
silica,  Si02  ,  or  as  silicates. 

Amorphous  silicon  is  formed  by  passing  vapor  of  SiCl4  over  heated  potassium; 
by  heating  magnesium  in  SiF4  vapor;  by  heating  a  mixture  of  Mg  and  Si02;  by 
electrolysis  of  a  fused  silicate.  It  is  readily  prepared  by  heating  a  mixture  of 
magnesium,  one  part,  with  sand,  four  parts,  in  a  wide  test-tube  of  hard  glass 
(Gattermann,  B.,  1889,  22,  186).  The  graphitoidal  form  is  crystalline  and  by 
many  is  said  to  'be  the  same  as  the  adamantine  form.  Method  of  preparation 
essentially  the  same  (Warren,  C.  N.,  1891,  63,  46).  The  crystalline  form  is  made 
by  fusing  a  silicate  or  K,SiF6  with  Al;  by  passing  vapors  of  SiCl4  over  heated 
Na  or  Al  in  a  carbon  crucible  (Deville,  A.  Ch.,  1857,  (3),  49,  62;  Devilie  and 
Caron,  A.  Ch.,  1863,  (3),  67,  435;  Woehler,  I.  c.). 

§249.   Silicon  dioxide.    SiO,  =  60.3. 
(Silicic  anhydride;  silica.) 

]  Silicic  acid.    H2Si03  =  78.316.          ^ 
SiIV0~22  and  H'2SiIV0-"r{ ,  0  =  Si  =  0  and  H  —  0  —  Si  —  0  —  H. 

1.  Properties. — Silica,  silicic  anhydride,  SiO2  ,  is  a  white,  stable,  infusible  solid; 
insoluble  in  water  or  acids;  soluble  in  fixed  alkalis  with  formation  of  silicates. 
Specific  iirurtty  of  quartz,  2.647  to  2.652;  of  amorphous  silica,  2.20  at  15.6°. 

Silicic  acid,  silicon  hydroxide,  H2SiO3  ,  is  a  white,  gelatinous  solid,  generally 
insoluble  in  water,  and  soluble  in  mineral  acids.  A  dilute  solution  in  water  is 
obtained  by  dialysis  of  the  fixed  alkali  silicate  with  an  excess  of  HC1  until 
the  chlorides  are  all  removed.  It  may  be  boiled  for  some  time  before  the  acid 
precipitates  out.  Upon  standing  silicic  acid  soon  separates. 

2.  Occurrence. — Silicon  is  never  found  free  in  nature;  it  is  always  combined 
with  oxyg'en  in  the  form  of  silicon  dioxide,  Si02  ,  as  quartz,  opal,  flint,  sand, 
etc.;  or  the  silicon  dioxide  is  in  combination  with  bases  as  silicates:  asbestos, 
soapstone,   mica,   cement,    glass,   etc.     All   geological   formations   except    chalk 
contain  silicon  as  the  dioxide  or  as  a  silicate. 

3.  Formation. — Crystalline  silica  is  formed  by  passing  silicon  fluoride  into 
water,   forming  silicic   acid  and  fluosilicic  acid:   3SiF4    -f-   3H2O   =  H2Si03    + 
2H,SiF6  .     The  precipitate  of  silicic  acid  is  dissolved  in  boiling  NaOH  and  then 
heated   in    sealed   tubes.     Below    180°    crystals    of   tridymite    are    formed,    and 
above  180°  crystals  of  quartz  (Maschke,  Pogff.,  1872,  145,  549). 

4.  [Preparation.— Pure   amorphous   silica   is   prepared   by   fusing   finely    pow- 
dered quartz  with  six  parts  of  sodium  Carbonate,  dissolving-  the  cooled  mass  in 
water,  and  pouring  into  fairly  concentrated  hydrochloric  acid.     The  precipitate 
is  filtered,  well  washed   and    ignited.     Or  SiF4    vapors  are  passed   into  wate.r 
(§246)   and  the  gelatinous  precipitate  washed,  dried  and  ignited.     Crystalline 


300  SILICON  DIOXIDE.  §249,  5. 

silica   is   prepared   by   fusing-   silicates   with    microcosmic    salt   or   with   borax 
(Rose,  /.  pr.,  1867,  101,  228). 

Silicic  acid. — The  various  hydroxides  of  silica  act  as  weak  acids.  Metasilicic 
acid,  HoSiOs  ,  has  been  isolated;  it  is  formed  by  decomposing  silicon  ethoxide, 
Si(OCaH5)4  ,  with  moist  air  (Ebelmen,  J.  pr.,  1846,  37,  359).  Also  by  dialysis  of 
a  mixture  of  sodium  silicate  with  an  excess  of  hydrochloric  acid  until  the 
chlorides  are  all  removed,  concentrating1,  allowing-  to  gelatinize,  and  drying 
over  sulphuric  acid.  Other  hydroxides,  acids,  have  been  isolated,  but  there  is 
some  question  as  to  their  exact  composition. 

5.  Solubilities.—  Silica,  Si02 ,  is  insoluble  in  water  or  acids  except  HF , 
which  dissolves  it  with  formation  of  gaseous  silicon  fluoride,  SiF4  (§246). 
Of  the  silicates  only  those  of  the  fixed  alkalis  are  soluble  in  water,  water 
glass.     These  silicates  in  solution  are  readily  decomposed  by  acids,  in- 
cluding carbonic  acid,  forming  silicic  acid,  gelatinous.     While  anhydrous 
silicic  anhydride,  Si02 ,  is  insoluble  in  mineral  acids,  the  freshly  precipi- 
tated hydroxide,  silicic  acid,  is   soluble  in  those   acids.     Silicic   acid  is 
decomposed  by  evaporation  to  dryness  in  presence  of  mineral  acids,  with 
separation  of  the  anhydrous  Si02  ;  which  is  insoluble  in  more  of  the  same 
acids,  which  previously  had  effected  its  solution. 

The  most  of  the  silicates  found  in  nature  are  of  complex  composition. 
They  are  combinations  of  Si02  with  bases.  They  are,  as  a  rule,  insoluble 
in  water  or  acids. 

6.  Reactions. — Solutions  of  the  alkali  silicates  precipitate  solutions  of 
all  other  metallic  salts  with  formation  of  insoluble  silicates;  they  are 
decomposed  by  acids  with  separation  of  silicic  acid,  a  gelatinous  precipi- 
tate, soluble  in  hydrochloric  acid.     Evaporation  decomposes  silicic  acid 
with  separation  of  insoluble  silicic  anhydride,  Si02 .     Ammonium  salts 
precipitate  gelatinous  silicic  acid  from  solutions  of  potassium  or  sodium 
silicate.     Therefore  in  the  process  of  analysis  the  silicic  acid,  not  removed 
in  the  first  group  by  hydrochloric .  acid,  will  be  precipitated  in  the  third 
group  on  the  addition  of  ammonium  chloride. 

Silica,  Si02 ,  is  soluble  in  hot  fixed  alkalis  forming  silicates ;  it  is  not 
soluble  in  ammonium  hydroxide,  nor  are  solutions  of  alkali  silicates  pre- 
cipitated on  addition  of  ammonium  hydroxide  as  they  are  on  the  addition 
of  ammonium  salts.  Boiling  Si02  with  the  fixed  alkali  carbonates  forms 
soluble  silicates  with  greater  or  less  readiness.  Nearly  all  silicates  are 
decomposed  by  heating  in  sealed  tubes  to  200°  with  concentrated  HC1  or 
H2S04 . 

7.  Ignition. — Silicates  fused  with  the  alkalis  form  soluble  alkali  sili- 
cates, and  oxides  of  the  metal  previously  in  combination.     If  alkali  car- 
bonates are  employed  the  same  products  are  formed  with  evolution  of 
C02 .     Preferably  a  mixture  (in  molecular  proportions)  of  potassium  and 
sodium  carbonates,  four  parts,  should  be  used  to  one  part  of  the  insoluble 
silicate.     Silica,  Si02 ,  is  also  changed  to  a  soluble  silicate  by  fusing  with 
fixed  alkali  hydroxides  or  carbonates. 


§250,  1.  PHOSPHORUS.  301 

Si02  does  not  react  with  K2S04  or  Na2S04  ,  even  when  fused  at  a  very  high 
temperature  (Mills  and  Mean\v<'ll,  ./.  ('..  is-si,  39,  f>:;:;).  In  the  fused  bead  of 
microcosm ic  salt  particles  of  silica  swim  iiiHliHxnlri'd.  If  a  silicate  be  taken, 
its  base  will,  in  most  cases,  be  dissolved  out,  leaving  a  "  skclclini  <>f  silica"  un- 
dissolved  in  the  liquid  bead.  But  with  a  head  of  sodium  carbonate,  silica  (and 
most  silicates)  fuse  to  a  clear  glass  of  silicate. 

Silica  is  separated  from  the  fixed  alkalis  in  natural  silicates,  by  mixing1  the 
latter  in  fine  powder  with  three  parts  of  precipitated  calcium  carbonate,  and 
one-half  part  of  ammonium  chloride,  and  healing  in  a  platinum  crucible  to 
redness  for  half  an  hour,  avoiding1  too  high  a  heat.  On  digesting  in  hot  water, 
the  solution  contains  all  the  alkali  metals,  as  chlorides,  with  calcium  chloride 
and  hydroxide. 

8.  Detection. — Silicates  are  detected  by  conversion  into  the  anhydride, 
Si02 .  The  silicate  is  fused  with  about  four  parts  of  a  mixture  of  potas- 
sium and  sodium  carbonates,  digested  with  warm  water,  filtered,  and 
evaporated  to  dryness  with  an  excess  of  hydrochloric  acid.  The  dry  resi- 
due is  moistened  with  concentrated  HC1  and  thoroughly  pulverized ;  water 
is  added  and  the  precipitate  of  Si02  is  thoroughly  washed.  Further  con- 
firmation may  be  obtained  by  warming  the  precipitate  of  SiO.,  with 
calcium  fluoride  and  sulphuric  acid  (in  lead  or  platinum  dishes),  forming 
the  gaseous  silicon  fluoride,  SiF4 .  This  is  passed  into  water  where  it  is 
decomposed  into  gelatinous  silicic  acid  and  fluosilicic  acid:  3SiF4  -f-  3H00 
=  H2SiO,  +  2H2SiF0  (§246).  Silica,  Si02 ,  is  usually  treated  as  directed 
for  silicates,  but  may  be  at  once  warmed  with  calcium  fluoride  and  sul- 
phuric acid. 

9.  Estimation. — The  compound  containing  a  silicate  or  silica  is  fused  with  fixed 
alkali  carbonates  as  directed  under  detection,  and  the  amount  of  well-washed  Si02 
determined  by  weighing  after  ignition. 

§250.  Phosphorus.     P  =  3i.04.     Usual  valence  three  or  five  (§11). 

1.  Properties. — Phosphorus  is  prepared  in  several  allotropic  modifications. 
Specific  gravity  of  the  yellow,  solid,  at  20°,  1.82321;  liquid,  at  40°,  1.74924; 
solid,  at  44.2°  1.814  (Damien,  1881).  At  ordinary  temperatures  it  is  brittle  and 
easily  pulverized.  At  44.1°  (Burgess,  Wash.  Acad.  of  Sci.,  1-18)  it  melts,  but 
may  be  cooled  in  some  instances  (under  an  alkaline  liquid)  as  low  as  +4°  with- 
out solidifying.  When  it  solidifies  from  these  lower  temperatures,  as  it  does 
by  stirring  with  a  solid  substance,  the  temperature  immediately  rises  to  about 
45°.  Boiling  point,  287.3°  at  762  mm.  pressure  (Schroetter,  A.,  1848,  68,  247; 
Kopp,  A.,  1855,  93,  129).  The  density  of  the  vapor  at  1040°  is  4.50  (Deville 
and  Troost,  C.  r.,  1863,  56,  891).  The  computed  density  for  the  molecule  P4 
is  4.294.  At  a  white  heat  the  density,  3.632,  indicates  dissociation  of  the  mole- 
cule to  P2  (Meyer  and  Biltz,  B.,  1889,  22,  725).  Specific  gravity  of  the  red  amor- 
phous modification  at  0°,  2.18  (Jolibois,  C.  r.,  149,  287-289). 

Ordinary  crystalline  yellow  stick  phosphorus  is  a  nearly  colorless,  trans- 
parent solid;  when  cooled  slowly  it  is  nearly  as  clear  as  water.  In  water  con- 
taining1 air  it  becomes  coated  with  a  thin  whitish  film.  If  melted  in  fairly 
large  quantities  and  cooled  slowly  it  forms  dodecahedral  and  octahedral  crys- 
tals (Whewell,  G.  N.,  1879,  39,  144).  Heated  in  absence  of  air  above  the  boiling 
point  it  sublimes  as  a.  colorless  gas,  depositing  lustrous  transparent  crystals 
(Blondlot,  C.  r.,  1866,  63,  397).  At  low  temperatures  phosphorus  oxidizes  slowly 
in  the  air  with  a  characteristic  odor,  probably  due  to  the  formation  of  ozone 
and  phosphorous  oxide,  P2O3  (Thorpe  and  Tutton,  ,7.  C.,  1890,  57,  573).  It  ignites 
spontaneously  in  the  air  at  44.5°,  burning  with  a  bright  yellowish  white  light 
producing  much  heat.  From  the  finely  divided  state,  as  from  the  evaporation 


302  PHOSPHORUS.  §250,  2. 

of  its  solution  in  carbon  disulphide,  it  ignites  spontaneously  at  temperatures 
at  which  the  compact  phosphorus  may  be  kept  for  days.  It  must  be  preserved 
under  water.  Great  precaution  should  be  taken  in  working  with  the  ordinary 
or  yellow  phosphorus.  Burns  caused  by  it  are  very  painful  and  heal  with 
great  difficulty.  Ordinary  phosphorus  is  luminous  in  the  dark,  but  it  has 
been  shown  that  the  presence  of  at  least  small  amounts  of  oxygen  are  neces- 
sary. The  presence  of  HoS  ,  SO,  ,  CS2  ,  Br  ,  Cl ,  etc.,  prevent  the  glowing 
(Schroetter,  J.  pr.,  1853,  58,  158;  Thorpe",  Nature,  1890,  41,  523).  Upon  heating 
in  absence  of  air,  better  in  sealed  tubes,  to  300°  it  is  changed  to  the  red  modi- 
fication (Meyer,  B.,  1882,  15,  297). 

Red  phosphorus  is  a  dull  carmine-red  tasteless  powder.  It  is  not  poisonous, 
while  the  ordinary  yellow  variety  is  intensely  poisonous,  200  to  500  milligrams 
being  sufficient  to  cause  death.  While  the  yellow  modification  is  so  readily 
and  dangerously  combustible  when  exposed  to  the  air  even  at  ordinary  tem- 
peratures, the  red  variety  needs  no  special  precautions  for  its  preservation. 
It  does  not  melt  when  heated  to  redness  in  sealed  tubes,  but  is  partially 
changed  to  the  yellow  crystalline  form  (Hittorf,  Pogg.,  1865,  126,  193).  If 
amorphous  phosphorus  be  distilled  in  the  absence  of  air,  it  is  changed  to  the 
( ly  -talline  form,  action  beginning  at  260°.  If  ordinary  red  phosphorus  is  heated  to 
4oJ°  for  sixty  hours,  in  absence  of  air,  it  becomes  red  pyromorphic  phosphorus 
wlio.se  density  is  2.37.  Heated  to  725°  this  phosphorus  melts,  and  if  suddenly 
cooled  becomes  a  violet-colored  mass  (Jolibois,  C.  r.,  149,  287-289).  Heated  in  the 
air  from  250°  to  260°  it  takes  fire  (Schroetter,  I.  c.).  A  black  crystalline  metallic 
variety  of  phosphorus  is  described  by  Hittorf  (/.  c.);  also  Remsen  and  Kaiser  (Am., 
1882,  <%  459)  describe  a  light  plastic  modification.  Phosphorus  is  largely  used 
in  match-making.  Yellow  phosphorus  is  used  in  the  ordinary  match,  and  the  red 
(amorphous)  in  the  safety  matches,  the  phosphorus  being  on  a  separate  surface. 

2.  Occurrence. — It  is  never  found  free  in  nature.     It  is  found  in  the  primitive 
rocks  as  calcium  phosphate,  occasionally  as  aluminum,  iron,  or  lead  phosphate, 
etc.     Plants  extract  it  from  the  soil,  and  animals  from  the  plants.     Hence  traces 
of  it  are  found  in  nearly  all  animal  and  vegetable  tissues;  more  abundantly 
in  the  seeds  of  plants  and  in  the  bones  of  animals. 

3.  Formation. — Ordinary  phosphorus  is  formed  by  heating  calcium  or  lead 
phosphates  with  charcoal.     The  yield  is  increased  by  mixing  the  charcoal  with 
sand  or  by  passing  HC1  gas  over  the  heated  mixture.     By  igniting  an  alkali 
or  alkaline  earth  phosphate  with  aluminum  (Rossel  and  Frank,  B..  1894,  27,  52). 
Red  phosphorus  is  formed  by  the  action  of  light,  heat  or  electricity  on  ordinary 
phosphorus  (Meyer,  B.,  1882,  15,  297).     By  heating  ordinary  phosphorus  with 
a  small  amount  of  iodine  (Brodie,  J.  pr.,  1853,  58,  171). 

4.  Preparation. — Ordinary   phosphorus    is   prepared    from    bones.     They    are 
first   burned,   which   leaves   a    residue,   consisting   chiefly   of   Ca3(P04)2;   then 
ELSO4    is    added,    producing   soluble    calcium    tetrahydrogen   diphosphate    (a). 
After  filtering  from  the  insoluble  calcium  sulphate  the  solution  is  evaporated 
and  ignited,  leaving  calcium  metaphosphate   (6).     Then  fused  with   charcoal, 
reducing  two-thirds  of  the  phosphorus  to  the  free  state  (c).     The  mixture  of 
sand,  SiO2  ,  with  the  charcoal  is  preferred,  in  which  case  the  whole  of  the 
phosphorus   is   reduced    (<?).     Hydrochloric   acid   passed    over   red-hot    calcium 
phosphate  and   charcoal  reduces  the  whole  of  the  phosphorus.     This  process 
works  well  in  the  laboratory,  and  has  also  been  successfully  employed  on  a 
larger  scale.     Either  of  the  calcium  phosphates  may  be  used  (e)  and  (f). 

(a)  Ca3(P04)2  +  2H2S04  =  2CaSO4  -f  CaH4(P04)2 

(6)  CaH4(PO4)2  +  ignition  —  Ca(P03),  +  2H2O 

(c)  3Ca(P03)2  +  IOC  =  Ca3(P04)2  +  10CO  +  P4 

(<?)  2Ca(P03)2  +  IOC  +  2Si02  =  2CaSiO3  +  P4  +  10CO 

(e)  2Ca3(P04)2  +  1GC  +  12HC1  =  GCaCl2  +  P4  +  16CO  -f  6H2 

(f)  2Ca(P03)2  +  12C  +  4HC1  =  2CaCL  +  P4  +  12CO  +  2H2 

Red  or  amorphous  phosphorus  is  prepared  by  heating  ordinary  phosphorus 
for  a  long  time  (40  hours)  at  240°  to  250°  in  absence  of  air.  At  260°  the  reverse 
change  takes  place.  If  the  heating  is  under  pressure  and  at  300°,  the  change 
to  the  red  phosphorus  is  almost  immediate.  It  is  washed  with  CS2  to  remove 
all  traces  of  yellow  phosphorus  and  is  dried  at  100°. 


§250,  ?,  PHOSPHORV8,  $03 

5.  Solubilities. — A  trace   of  phosphorus   dissolves  in  water.     Alcohol 
dissolves  0.4,  ether  0.9,  olive  oil  1.0,  and  turpentine  2.5  per  cent  of  it, 
while  carbon  disulphide  dissolves  10  to  15  times  its  own  weight.     lied 
phosphorus  is  insoluble  in  water,  ether,  or  carbon  disulphide. 

6.  Reactions. — When  phosphorus  is  boiled  with  a  fixed  alkali  or  alkaline 
earth    hydroxide,   phosphorus   hydride,   phosphine    (§249),    PH3 ,    and   a 
hypophosphite  '(§250)    are   formed.     Phosphorus,    when   warmed   in    an 
atmosphere  of  N  or  C02 ,  combines  directly  with  many  metals  to  form 
phosphides.     These   phosphides   are   usually   brittle   solids   decomposing 
with  water  or  dilute  acids  with  formation  of  phosphoretted  hydrogen, 
PH3 .     In  nearly  all  the  reactions  of  phosphorus  both  varieties  react  the 
same,  the  red  variety  with  much  less  intensity,  and  frequently  requiring 
the  aid  of  heat.     It  is  ignited  when  brought  in  contact  with  Pb02 ,  Pb304 
HgO ,  Ag20 ,  CrO;i ,  K2Cr207  and  when  heated  with  CuO  or  Mn02 .     Solu- 
tions of  platinum,  gold,  silver,  and  copper  salts  are  decomposed  by  phos- 
phorus with  separation  of  the  corresponding  metal  (Boettger,  J.  C.,  1874, 
27,  1060). 

With  HN03,  H3P04  and  NO  are  formed;  when  heated  with  KN03  a 
rapid  oxidation  takes  place. 

It  combines  with  oxygen,  forming  P203  or  P20.  .  With  yellow  phos- 
phorus the  reaction  begins  at  ordinary  temperature;  with  the  red  variety 
not  till  heated  to  250°  to  260°  (Baker,  J.  (7.,  1885,  47,  349). 

Water  is  decomposed  at  250°,  forming  PH.,  and  H:!P04  (Schroetter,  I.  c.}. 

Combination  with  red  phosphorus  and  sulphur  takes  place  at  ordinary 
temperatures,  forming  P2S3  or  P2S5 ,  depending  upon  the  proportion  of 
each  employed  (Kekule,  A.,  1854,  90,  310).  With  ordinary  phosphorus 
the  action  is  explosive.  Sulphides  of  phosphorus  with  formulas  P4S3 
and  P4S7  have  also  been  prepared.  (Stock,  B.,  43,  414,  1223.) 

Cl  or  Br  react  with  incandescence  at  ordinary  temperatures,  forming 
trihalogen  or  pentahalogen  compounds,  depending  upon  the  amount  of 
halogen  employed.  With  iodine,  PI3  is  formed. 

The  halogen  compounds  of  phosphorus  are  decomposed  by  water  with 
formation  of  the  corresponding  hydracids  and  phosphorous  or  phosphoric 
acids,  depending  upon  the  degree  of  oxidation  of  the  phosphorus.  In 
the  presence  of  water  phosphorus  is  oxidized  to  H3P04  by  Cl ,  Br ,  I . 
HC103 ,  HBr03 ,  or  HI03  with  formation  of  the  corresponding  hydracid: 
P4  +  10C12  +  16H20  ==  4H3P04  +  20HC1 . 

7.  Ignition. — When  sodium  carbonate  is  heated  to  redness  with  phosphorus, 
the  carbonic  anhydride  is  reduced  and  carbon  is  set  free.     Phosphorus  heated 
with   magnesium  in   a   vapor  of  carbon  dioxide   forms  P2Mg3  ,   which   can    be 
heated  to  redness  in  absence  of  air  without  decomposition.     Heated  in  the  air 
it  becomes  oxidized  (Blunt,  A.  Ch.,  1865,  (4),  5,  487).     Phosphorus  also  combines 
with  Cu  ,  Ag  ,  Cd  ,  Zn  and  Sn  when  it  is  heated  with  these  elements  in  sealed 
tubes.     It  does  not  combine  with  Al  and  but  slightly  with  Fe   (Emmerling-, 
J.  C.,  1879,  36,  508), 


304  PHOKPHINE— HYPOPHOSPHOROUS  ACID.  §250,  8. 

8.  Detection. — By  its  phosphorescence;  by  formation  of  PH3  when 
boiled  with  KOH  (Hofmann,  .#.,  1871,  4,  200);  by  oxidation  to  H.PC^  and 
detection  as  such  (§75,  6d). 

9.  Estimation. — Oxidation  to  H3P04  ,  precipitation  with  magnesia  mixture  as 
Mg-NH4P04  ,  ignition  to,  and  weighing  as  Mg2P2OT   (§189,  9). 


$251.  Phosphine.     PH3=  34.064. 
P-'"H'3,H  —  P  —  H. 


Phosphine,  PH3  ,  is  a  colorless  gas  having  a  very  disagreeable  odor.  As 
usually  prepared,  it  is  spontaneously  inflammable,  burning  in  the  air  with 
formation  of  metaphosphoric  acid:  2PH3  +  402  =  2HP03  +  2H20  .  It  is 
liquified  and  frozen  at  very  low  temperatures;  boiling  point,  about  — 85°; 
melting  point,  — 132.5°  (Olszewski,  M.,  1886,  7,  371).  It  is  very  poisonous,  spar- 
ingly soluble  in  water,  which  solution  has  the  peculiar  odor  of  the  gas  and  has 
an  exceedingly  bitter  taste.  It  is  formed  by  boiling  phosphorus  with  a  fixed 
alkali  or  alkaline  earth  hydroxide  (a);  by  ignition  of  H3PO2  or  H3PO3  (6);  by 
ignition  of  hypophosphites  (c) ;  by  the  decomposition  of  the  alkaline  earth 
phosphides  with  water  or  dilute  acids  (d) : 

(a)     P4  +  3KOH  +  3H20  =  3KH2P02  +  PH3 

(6)     2H3P02  =  HP03  +  PH3  +  H20 
4H3P03  =  3HP03  +  PH3  +  3H2O 

(c)  4NaH2P02  =  Na4P207  +  2PH3  +  H20 

(d)  Ca3P2  +  GH20  =  3Ca(OH)2  +  2PH3 
Ca3P2  +  6HC1  =  3CaCl2  +  2PH3 

It  is  a  strong  reducing  agent;  transposes  many  metallic  solutions:  3CuS04  -f- 
2PH3  =  Cu3P2  +  3H,SO4;  reduces  solutions  of  silver  and  gold  to  the  metallic 
state:  8AgNO3  +  PH3"  +  4H2O  =  H3P04  +  8HNO3  +  8Ag;  is  oxidized  to  H3PO4 
by  hot  H,S04  ,  Cl  ,  HC1O  ,  HNO2  ,  HN03  ,  H:,AsO4  ,  etc.  A  liquid  phosphorus 
hydride,  P2H4  ,  and  a  solid,  P4H2  ,  are  known  (Besson,  C.  r.,  1890,  111,  972; 
Gattermann  and  Hausknecht,  B.,  1890,  23,  1174). 


§252.  Hypophosphorous  acid.     H3P02  =  66.064  . 

H 

I 
H'3PO-"2.     H  — 0  — P  =  0. 

I 

H 

1.  Properties. — Hypophosphorous  acid  was  discovered  in  1816  by  Dulong  (A.  C7i., 
1816,  2,  141).     It  is  a  colorless  syrupy  liquid;  specific  gravity,  1.493  at  18.8°.     At 
17.4°  it  becomes  a  white  crystalline  solid  (Thomsen,  B.,  1874,  7,  994).     Although 
containing  three  hydrogen  atoms  it  forms  but  one  series  of  salts,  e.  g.,  NaH2PO2, 
Ba(H2PO.)2  ,  etc. 

2.  Occurrence. — Not  found  in  nature. 

3.  Formation. — All  ordinary  metols  form  hypophosphites  except  tin,  copper 
and  mercurosum.     Silver  and  ferric  hypophosphites  are  very  unstable.     (1)  A 


£252,  8,  iiYJ'oriioxi'noh'ucti  ACID.  305 

few  metals,  such  as  zinc  and  iron,  dissolve  in  H3PO.,  ,  giving  off  hydrogen  and 
forming  a  hypophosphite.  (:.')  The  alkali  and  alkaline  earth  salts  may  be 
formed  by  boiling  phosphorus  with  the  hydroxides  (Maw row  and  Muthmann, 
Z.  anycic.,  1896,  ii,  268).  (3)  As  all  bypophosphitefi  are  soluble,  none  can  be 
formed  by  precipitation.  All  may  be  formed  from  their  sulphates  by  trans- 
position with  barium  hypophosphite.  (4)  All  may  be  made  by  adding"  H8P02 
to  the  carbonate's  or  hxdroxides  of  the  metals. 

4.  Preparation. — To  prepare  pure  H3PO2  ,  BaO  and  P   (in  small  pieces)  are 
warmed    in    an    open    dish    with    water   until    PH3    ceases    to    be    evolved.     The 
liquid  is  filtered  and  excess  of  BaO  is  removed   by  passing  in  CO.  .     After  again 
filtering,  the  liquid  is  evaporated  to  crystallization  of  the  barium   salt.     This 
is  dissolved   in  water  and   decomposed    by   the   calculated   quantity   of   H2SO4  . 
The  solution  is  filtered  and  evaporated  in  an  open  dish,  care  being  taken  not  to 
heat  above  110°  .     l.'pon  cooling  the  white  crystalline  tablets  are  obtained. 

5.  Solubilities. — The  free  acid  is  readily  miscible  in  water  in  all  proportions. 
The  salts  are  all  soluble  in  water,  a  number  of  them  are  soluble  in  alcohol. 

6.  .Reactions. — A. — With  metals   and   their    compounds.       Hypophosphorous 
acid  is  a  very  powerful  reducing  agent,  being  oxidized  to  phosphoric  acid  or  a 
phosphate. 

1.  Pbiv  becomes  Pb"  in  acid  or  alkaline  mixture. 

2.  Ag'  becomes  Ag-°  in  aeid  or  alkaline  mixture. 

3.  Hg"  becomes  Hg'  and  then  Hg-°  in  acid  or  alkaline  mixture. 
4-  Asv  and  As'"  become  As0  in  presence  of  HC1 . 

5.  Bi'"  becomes  Bi°  in  presence  of  alkalis  or.  acetic  acid. 

6.  Cu"  becomes  Cu2H2  and  on  boiling  Cu°  (separation  from  Cd). 

7.  Fe"'  becomes  Fe"  ,  no  action  in  alkaline  mixture. 

8.  CrVi  becomes  Or'"  ,  no  action  in  alkaline  mixture. 

9.  Co'"  becomes  Co"  ,  no  action  in  alkaline  mixture. 

10.  Ni'"  becomes  Ni"  ,  no  action  in  alkaline  mixture.^ 

11.  Mn"+n  becomes  Mn"  in  acid  solution. 

12.  Mniv+n  becomes  Mniv  in  alkaline  mixture. 

B. — With  non-metals  and  their  compounds. 

1.  H,Fe(CN)G  becomes  H.,Fe(CN)6  . 

2.  HN03  and  HNO2  become  WO  . 

3.  H3PO_,  on  heating  becomes  H3P04  and  PH3  . 

4-  H,S03  becomes  free  sulphur  with  formation  of  some  H2S  (Ponndorf,  J.  0., 

1877,31,  275). 
H,SO4  becomes  first  H,,SO3  then  S  .     See  above. 

5.  Cl  becomes  HC1  in  acid  mixture,  a  chloride  with  alkalis. 
HC1O  and  HC1O3  form  same  products  as  Cl  . 

6.  Br  becomes  HBr  in  acid  mixture,  a  bromide  with  alkalis. 
HBr03  forms  HBr  . 

7.  I  forms  HI  ,  in  alkaline  mixtures  an  iodide. 

HI,  dry,  reacts  violently,  forming  H3PO3  and  PH4I  (Ponndorf,  I.e.). 
HIO8  forms  HI  . 

7.  Ignition. — On  ignition  hypophosphites  leave  pyrophosphates,  evolving  PE8. 
The  acid  decomposes  on  heating  to  PH3  and  H3PO'4  (or  HPO3  if  at  a  red  heat), 

8.  Detection.— Hypophosphorous  acid  may  be  known  from  phosphorous 
acid  by  adding  cupric  sulphate  to  the  free  acid  and  heating  the  solution 
to  55°.  With  hypophosphorous  acid  a  reddish-black  precipitate  of  copper 
hydride  (Cu2H2)  is  thrown  down,  which,  when  heated  in  the  liquid  to  100°, 
is  decomposed  with  the  deposition  of  the  metal  and  the  evolution  of 
hydrogen,  whilst  with  phosphorous  acid  the  metal  is  precipitated  and 
hydrogen  evolved,  but  no  Cu2H2  is  formed.  Further,  hypophosphorous 
acid  reduces  the  permanganates  immediately,  but  phosphorous  acid  only 
after  some  time.  Phosphites  precipitate  barium,  strontium,  and  calcium 


306  PHOSPHOROUS  ACID.  §252,  9. 

salts,  while  hypophosphites  do  not.  When  hypophosphorous  acid  is 
treated  with  zinc  and  sulphuric  acid  it  is  converted  into  phosphoretted 
hydrogen.  On  boiling  hypophosphorous  acid  with  excess  of  alkali  hydrox- 
ide, first  a  phosphite  then  a  phosphate  is  formed,  with  evolution  of 
hydrogen. 

9.  Estimation. —   (1)   By  oxidation  with  nitric  acid  and  then  proceeding  as 

with   phosphoric   acid.     (2)    By    mercuric   chloride    acidulated    with    HC1;    the 

temperature    must    not   rise    above    60°,    otherwise    metallic    mercury    will    be 

formed.     The  precipitated  Hg-Cl ,  after  washing-  and  drying-  at  100°,  is  Aveighed. 

NaH,P02  +  4HgCL  +  2H,0  =  4HgCl  +  H3P04  +  NaCl  +  3HC1 


j253.  Phosphorous  acid.     H,P03  ==  82.064  . 

H 

I 
H'3P"0-"3  ,H  —  0  —  P  —  0  —  H. 


1.  Properties. — Phosphorous  anhydride,  P203  ,  is  a  snow-white  solid,  melting 
at  22.5°,  and  boiling  at  173.1°    (Thorpe  and  Tutton,  J.  C.,  1890,  57,  f>45).     The 
vapor  density  of  the  gaseous  oxide  indicates  the  molecule  to  be  P406  .     tiitccific 
yntvUy  of  the  liquid  at  21°,  1.9431;  of  the  solid  at  the  same  temperature,  2.135. 
It  has  an  odor  resembling  that  of  phosphorus.     Heated   in   a   sealed  tube   at 
200°  it  decomposes  into  P2O4  and  P  (T.  and  T.,  J.  C.,  1891,  59,  1019).     It  reacts 
slowly  with  cold  \vater,  forming  H3PO3;  with  hot  water  the  reaction  is  violent 
and  PH3  is  evolved.     Upon  exposure  to  the  air  it  oxidizes  to  P^OG  . 

The  acid,  H3P03  ,  is  a  crystalline  solid,  very  deliquescent,  melting  at  74° 
(Hurtzig  and  Geuther,  A.,  1859,  111,  171).  It  is  a  dibasic  arid,  forming  no 
tribasic  salts  (Amat,  (\  r.,  1889,  108,  403).  One  or  two  of  the  hydrogen  atoms 
are  replaceable  by  metals  forming  acid  or  normal  salts.  The  third  hydrogen 
is  never  replaced  by  a  metal,  but  may  be  replaced  by  organic  radicles  (Railton, 
J.  C.,  1855,  7,  216;  Michaelis,  J.  0.,  1875,  28,  1160).  Neither  meta  nor  pyro- 
phosphorous  acids  are  known,  but  a  number  of  pyrophosphites  have  been  pre- 
pared (Amat,  C.  r.,  1888,  106,  1400;  1889,  108,  1056;  1890,  110,  1191  and  901; 
A.  Ch.,  1891,  (6),  24,  289). 

2.  Occurrence. — Does  not  occur  in  nature. 

3.  Formation. — P2O3    is    formed    together    with    P205    when    phosphorus    is 
ignited  in  the  air.     H3PO3  is  formed  together  with  H3P04   when  phosphorus 
is  oxidized  with  HNO3;  by  the  oxidation  of  H3P02:  by  the  action  of  P  upon  a 
concentrated  solution  of  CuS04   in   absence  of  air:   3CuS04    +   P4    +   6ELO  = 
Cu3P2  +  2H3PO3  +  3H2SO4   (Schiff,  A.,  1860,  114,  200). 

4.  Preparation — To    prepare    phosphorous    anhydride,    P,O,  ,    phosphorus    is 
burned  in  a  tube  with  an  insufficient  supply  of  air   (Thorpe  and  Tutton,  I.e.). 
The  acid,  H3P03  ,  is  prepared  by  dissolving  the  anhydride  in  cold  water;   by 
decomposing'  PC13  with  water  (Hurtzig  and  Geuther,  I.  c.}. 

r>.  Solubilities. — The  acid  is  miscible  in  water  in  all  proportions.  Alkali 
phosphites  are  soluble  in  water,  most  others  are  insoluble  (distinction  from 
hypophosphites) . 

6.  Reactions.— Phosphorous  acid  is  a  strong  reducing  agent,  oxidizing  to 
phosphoric  acid  when  exposed  to  the  air.  It  reduces  salts  of  silver  and  gold  to 
the  metallic  state  and  is  changed  to  phosphoric  acid  by  most  of  the  strong 
oxidizing  acids  and  by  many  of  the  higher  metallic  oxides.  HgCL  becomes 
Hg-Cl  and  then  Hg°  ,  CuCL  becomes  CuCl  then  Cu°  (Rammelsberg,  J.  C.,  187;t, 


§255,  1.  H7POPHOSPHORIC  ACID— PHOSPHORIC  ACID.  307 

26,  13).  Concentrated  H2S04  with  heat  forms  H,P04  and  S02  (Adie,  J.  C.,  1891, 
59,  230).  H,S03  forms  H2S  and  H3P04  (Woehler,  A.,  1841,  39,  252).  Nascent 
hydrogen  (Zn  and  H,S04)  produce  PH3  (Dusart,  C.  r.,  1856,  43,  11  :>(>). 

7.  Ignition.— The  acid  is  decomposed  by  ignition,  forming  HPO;i   and   P  or 
PH,   (Yigier,  Bl,  1869,  (2),  11,  125;  Hurtzig  and  Geuther,  7.  r.).     Phosphites  Jin- 
decomposed   by  heat,  leaving  a  pyrophosphate  and   a    phosphide  and  evolving 
PH3  or  H  (Rammelsberg1,  /?.,  1876,9,  1577;  and  Kraut,  A.,  1875,  177,  274). 

8.  Detection. — By   oxidation  to   H3P04    and   detection    as   such.     Tt   is   distin- 
guished   from    hypophosphorons    acid    by    reducing   CuSOt    to    Cu°,    while    the 
latter  forms  Cu2H,.;  also  by  the  solubilities  of  the  salts  (§252,  8).     Its  reactions 
with   oxidizing  agents   distinguish   it  with   hypophosphorous   acid   from   phos- 
phoric acid. 

9.  Estimation. — By  oxidation  to  H3PO4  and  estimation  as  such. 


§254.  Hypophosphoric  acid.     H4P206  =  162.112. 

0       0 
II       II 
H'4PIV20-"6,H  —  0  —  P  —  P  —  0  —  H. 

I        I 
0       0 

I      I 

H      H 

Hypophosphoric  acid  is  formed  together  with  phosphorous  and  phosphoric 
acids  by  slowly  oxidizing  phosphorus  in  moist  air  (Salzer,  A.,  1885,  232,  1 1  1 
and  271);  also  by  oxidizing  phosphorus  with  dilute  HN03  in  presence  of  silver 
nitrate  (Philipp,  B.,  1885,  18,  749).  It  consists  of  small  colorless  hygroscopic 
crystals  which  melt  at  55°.  It  decomposes  when  heated  to  70°  into  H3PO3  and 
HPO3  ,  and  at  120°  gives  H4P207  and  PH3  (Joly,  C.  r.,  1886,  102,  110  and  760). 
It  is  oxidized  to  H3PO4  by  warm  HNO,  ,  slowly  by  KMn04  in  the  cold,  rapidly 
when  heated.  It  is  not  oxidized  by  H202  ,  chlorine  water  or  H2CrO4;  HgCl, 
becomes  HgCl  (Amat,  C.  r.,  1890,  111,  676).  It  is  not  reduced  by  Zn  and  H2SO, 
(distinction  from  H3P02  and  H3P03).  With  a  solution  of  silver  nitrate  it  gives 
a  white  precipitate  which  does  not  blacken  in  the  light  (distinction  from  H3P02 
and  H3P03).  It  forms  four  series  of  salts,  all  four  hydrogen  atoms  being 
replaceable  by  a  metal.  The  hypophosphates  are  much  more  stable  towards 
oxidizing  agents  than  hypophosphites  or  phosphites. 


§255.  Phosphoric  acid.     H3P04  =  98.064  . 

0 
II 

H'3PvO-"4,  H  — 0  — P  — 0  — H. 
I 

0 

I 

H 

1.  Properties. — Phosphoric  anhydride,  P2O5  *,  is  a  white,  flakey,  very  delique- 
scent solid,  fusible,  subliming  undecomposed  at  a  red  heat.  It  is  very  soluble 
in  water,  forming  three  varieties  of  phosphoric  acid:  ortlto,  H3P04;  meta,  HPO^; 

*  According  to  Tildeii  and  Barnett  (J".  C.,  1896,  69,  154)  the  molecule  is  P4O10  not  P2O5 ;  P4Oa 
not  PUO3 1  Thorpe  and  Tutton,  J.  C.,  1891,  59, 1032) ;  ana  P4S6  not  PaS3  (Isambert,  C.r.,  1886, 1O», 
1386). 


308  PHOSPHORIC  ACID.  §255,  2. 

and  pyro,  H4P2OT  .  Orthophosphoric  acid  is  a  translucent,  feebly  crystallizable 
and  very  deliquescent  soft  solid.  Specific  t/niriti/,  1.88  (Schiff,  A.,  1860,  113,  183); 
melting  point,  41.75°  (Berthelot,  BL,  1878,  (2), '29,  3).  It  is  changed  by  heat, 
first  to  pyrophosphoric  acid,  then  to  metaphosphoric  acid.  Orthophosphoric 
acid  forms  three  classes  of  salts:  M'H2PO4  ,  primary,  monobasic  or  mono- 
metallic phosphates:  M'2HPO4  ,  secondary,  dibasic  or  dimetallic  phosphates; 
and  M'3P04  ,  tertiary,  tribasic,  trimetallic  or  normal  phosphates.  The  first 
two  are  acid  salts,  but  Na2H?04  is  alkaline  to  test  paper.  Metaphosphoric 
acid,  HPO3  ,  H  —  O  —  P  —  O  ,  is  a  white  waxy  solid,  volatile  at  a  red  heat 
II 
O 

(ordinary  glacial  phosphoric  acid  owes  its  hardness  to  the  universal  presence  of 
sodium  metaphosphate).  It  is  a  monobasic  acid,  but  there  are  various  poly- 
meric modifications,  distinguished  from  each  other  chiefly  by  physical  differ- 
ences of  the  acids  and  their  salts  (Tammann,  Z.  pliys.  Ch.,  1890,  6,  122). 

O  0 

II  II 

Pyrophosphoric    acid,    H4P2O7  ,    H—  O  —  P  —  0— P  —  O  —  H,isa  glass-like 

I  l 

O  O 

I  I 

H  H 

solid  (Peligot,  A.  Ch.,  1840,  (2),  73,  286),  very  soluble  in,  but  unchanged  by, 
water  at  ordinary  temperature;  changed  by  boiling  water  to  H3PO4  .  Heated 
to  redness  HPO3  is  formed.  It  forms  two  classes  of  salts:  M'oH->PoO7  and 
M'4P207  . 

2.  Occurrence. — Phosphates  of  Al ,  Ca  ,  Mg  and  Pb  are  widely  distributed  in 
minerals.     Guano  consists   quite   largely   of   calcium   phosphate.     Calcium   and 
magnesium  phosphates  are  found  in  the  bones  of  animals  and  in  the  ashes  of 
plants.     The  free  acids  are  not  found  in  nature. 

3.  Formation. — Phosphoric  a  nil  yd  rule,  P.O.-,  ,  is  formed  by  burning  phosphorus 
in  great  excess  of  air;   also  by   burning  phosphorus  in  NO  ,   N02  ,   or   C102  . 
Orthophosphoric  acid,   H3P04  ,   is   formed   by   long   exposure   of   phosphorus   to 
moist  air,  or  by  oxidation  with  HN03;  by  oxidation  of  HSPO2  or  H3POS  with 
the  halogens,  HN03  ,  HC103  ,  etc.;  by  treating  P20r>  ,  HP03  ,  or  H4P2OT   with 
boiling  water;  by  combustion  of  PH3  in  moist  air;  and  by  action  of  water  on 
PC15  .     It  is  also  formed  from  metallic  phosphates  by  transposition  with  acids 
in  cases  where  a  precipitate  results,  as  a  lead  or  barium  phosphate  with  sul- 
phuric acid,  or  silver  phosphate  with  hydrochloric  acid.     But  wrhen  the  pro- 
ducts are  all  soluble,  as  calcium  phosphate  with  acetic  acid  or  sodium  phosphate 
with  sulphuric  acid,  the  transposition  is  only  partial;  so  that  unmixed  phos- 
phoric acid  is  not  obtained.     A  non-volatile  acid,  like  phosphoric,  is  not  sepa- 
rated from  liquid  mixtures,  as  the  volatile  acids  are,  like  hydrochloric.     The 
change  represented  by  equation    («)   can  be  rerifiecl,  that  is,  pure  phosphoric 
acid  can  be  separated;  but  the  changes  shown  in  equations  (ft)  and  (c)  do  not 
comprise  the  whole  of  the  material  taken.     In  the  operation   (&)   some  sodium 
phosphate   and    some   nitric    acid   will   be   left,    and   in    (c)    some    trihydrogen 
phosphate  will  no  doubt  be  made. 

a.  CaH4(P04)2  +  H2C204  =  CaC,04  +  2H3P04 

6.  Na2HP04  +  2HNO3  =  2NaNO8  +  H3PO4 

and  Na2HPO4  -f  HNO3  =  NaNO,  +  NaH2PO4 

c.  2CaHP04  +  2HC1  =  CaCL  +  CaH4(P04)2 

Metaphosphoric  odd  is  formed  by  treating  P20.,  with  cold  water;  by  decom- 
position of  lead  metaphosphate  with  H2S  or  of  the  barium  salt  with  H2SO4; 
by  ignition  to  dull  redness  of  phosphorus  or  any  of  its  acids  in  the  presence 
of  air  and  moisture. 

Pyrophosphoric  acid,  H4P207  ,  is  formed  by  the  decomposition  of  lead  pyro- 
phosphate,  Pb,P2O7  ,  with  H2S  or  of  the'  corresponding  barium  salt  with 
H2S04;  or  by  heating  H3PO4  to  a  little  above  200°  until  no  yellow  silver 
phosphate,  Ag1:,POl  ,  is  obtained  on  dissolving-  in  water  and  Ireatment  with 
silver  nitrate  after  neutralization  with  NH4OH  . 


§255,  5.  i'HOMi>noitir  ACID.  309 

4.  Preparation. — To  prepare  P205 ,  phosphorus  is  burned  in  a  slow  cur- 
rent of  dry  oiygen  healing  to  about  300°,  then  in  a  more  rapid  current 
of  the  gas,  and  iinally  the  P,0-  is  distilled  in  an  atmosphere  of  oxygen 
(Shenstone,   \Vnll*'  Die.,   1894,  IV,  141).     H3P04  is  prepared  by  warming 
phosphorus,  one  part,  with  nitric  acid,  */;.  gr.  1.20,  ten  to  twelve  parts, 
with  addition  of  300  to  600  milligrams  of  iodine  to  100  grams  of  phos- 
phorus,  until   the   phosphorus   is   completely   dissolved.     The   excess   of 
HNO.j  is  removed  by  evaporation,  water  is  added  and  the  solution  is  sat- 
urated with  H2S  to  remove  any  arsenic  that  may  be  present.     The  solution 
is  then  evaporated  to  a  syrupy  consistency  at  temperatures  not  above 
150°   (Krauthausen,  Arch.  Pharm.,  1877,  210,  410;  Huskisson,  B.,  1884, 
17,  161).     Many  orthophosphates  are  formed  by  the  action  of  H3P04  upon 
metallic  oxides  or  carbonates ;  by  the  reaction  between  an  alkali  phosphate 
and  a  soluble  salt  of  the  heavy  metal;  by  fusion  of  a  metaphosphate  with 
the  corresponding  metallic  oxide  or  hydroxide;  also  by  long  continued 
boiling  of  meta  or  pyrophosphates.     Metaphosphates  are  formed  by  double 
decomposition  with  NaPO,  or  by  fusion  of  a  monobasic  phosphate  or  any 
phosphate  having  but  one  hydrogen  equivalent  substituted  for  a  metal, 
the  oxide  of  which  is  non-volatile,  e.  g.,  NaNH4HP04 .     Pyrophosphates 
are  formed  by  double  decomposition  with  Na4P20T  ;  by  action  of  H4P207 
on  certain  oxides  or  hydroxides;  also  by  ignition  of  dibasic  orthophos- 
phates, e.  g.,  Na2HP04 .     Na,H,P207  may  be  prepared  by  titrating  a  sat- 
urated solution  of  Na4P20T  with  HN03  until  the  solution  gives  a  red  color 
with  methyl  orange.     Upon  standing  the  salt  separates  in  large  crystals 
(Knorre,  Z.  angew.,  1892,  639). 

5.  Solubilities. — All  the  phosphoric  acids  are  readily  soluble  in  water, 
as  are  all  alkali  phosphates.     Alkali  primary  orthophosphates  have  an 
acid  reaction  in  their  solutions;  alkali  secondary  and  tertiary  phosphates 
are  alkaline  in  their  solutions;  the  latter  is  easily  decomposed,  even  by 
C02 ,  forming  the  secondary  salt.     A  number  of  non-alkali  primary  ortho- 
phosphates  are  soluble  in  water,  e.  g.9  CaH4(P04)2 .     All  normal  and  di- 
metallic  orthophosphates  are  insoluble  except  those  of  the  alkalis.     The 
normal  and  dimetallic  phosphates  of  the  alkalis  precipitate  solutions  of 
all  other  salts.     The  precipitate  is  a  normal,  dimetallic,  or  basic  phos- 
phate, except  that  with  the  chlorides  of  mercury  and  antimony  it  is  not 
a  phosphate  but  an  oxide  or  an  oxychloride. 

All  phosphates  are  dissolved  or  transposed  by  HNO., ,  HC1 ,  or  H2S04 , 
and  all  are  dissolved  by  HC2H,02  except  those  of  Pb ,  Al  and  Fe'"  .  All 
are  soluble  in  H3P04  except  those  of  lead,  tin,  mercury,  and  bismuth. 

The  non-alkali  meta  and  pyrophosphates  are  generally  insoluble  in 
water.  The  pyrophosphates  of  the  alkaline  earth  metals  are  difficultly  solu- 
ble in  acetic  acid.  The  most  of  the  pyrophosphates  of  the  heavy  metals, 


:n<>  PHOSPHORIC  ACID.  §255,  6^4. 

except  silver,  are  soluble  in  solutions  of  alkali  pyrophosphates,  as  double 
pyro phosphates  soluble  in  water  (distinction  from  orthophosphates).  Ferric 
iron  as  a  double  pyre-phosphate  loses  the  characteristic  properties  of  that 
metal  (Persoz,  J.  C.,  1849,  1,  183).  Phosphates  are  insoluble  in  alcohol. 

0.  Reactions. — A. — With  metals  and  their  compounds. — Phosphoric  acid  dis- 
solves some  rnetals,  e.  g.,  Fe  ,  Zn  and  Mg  with  evolution  of  hydrogen.  It  unites 
with  the'  oxides  and  hydroxides  of  the  alkalis  and  alkaline  earths  and  with 
other  freshly  precipitated  oxides  and  hydroxides  except  perhaps  antimonous 
oxide.  It  also  decomposes  all  carbonates  evolving-  CO2  .  Phosphates  are  formed 
in  the  above  reactions,  the  composition  of  which  depends  upon  the  conditions 
of  the  experiment. 

Free  orthophosphoric  acid  is  not  precipitated  by  ordinary  salts  of  third, 
fourth  and  fifth  group  metals  (in  instance  of  ferric  chloride,  a  distinction  from 
pyrophosphoric  acid  and  metaphosphoric  acid),*  but  is  precipitated  in  part  b}r 
silver  nitrate,  and  lead  nitrate  and  acetate.  Ammoniacal  solution  of  calcium 
chloride  or  of  barium  chloride  precipitates  the  normal  phosphate. 

Free  metaphosphoric  acid  precipitates  solutions  of  silver  nitrate,  lead  nitrate, 
and  lead  acetate,  the  precipitates  being1  insoluble  in  excess  of  metaphosphoric 
acid,  and  soluble  in  moderately  dilute  nitric  acid.  Barium,  calcium  and  ferrous 
chlorides,  and  magnesium,  aluminum,  and  ferrous  sulphates,  are  not  precipi- 
tated by  free  metaphosphoric  acid.  Ferric  chloride  is  precipitated,  a  distinc- 
tion from  orthophosphoric  acid. 

Free  pyrophosphoric  acid  gives  precipitates  with  solutions  of  silver  nitrate, 
lead  nitrate  or  acetate,  and  ferric  chloride;  no  precipitates  with  barium  or 
calcium  chloride,  or  writh  magnesium  or  ferrous  sulphate. 

Orthophosphoric  acid — or  an  orthophosphate  with  acetic  acid — does  not  coagu- 
late egg  albumen  or  gelatine.  This  is  a  distinction  of  both  orthophosphoric 
acid  and  pyrophosphoric  acid  from  metaphosphoric  acid. 

With  silver  nitrate  soluble  orthophosphates  in  neutral  solution  form 
silver  orthophosphate,  Ag3P04 ,  yellow ;  with  metaphosphates,  silver  m eta- 
phosphate,  AgP03 ,  white ;  and  with  pyrophosphates,  silver  pyrophosphate, 
Ag4P207 ,  white,  all  soluble  in  ammonium  hydroxide.  Silver  metaphos- 
phate  is  soluble  in  excess  of  an  alkali  metaphosphate  (distinction  from 
pyrophosphates). 

If  a  disodium  or  dipotassium  orthophosphate  is  added  to  solution  of  silver 
nitrate,  free  acid  is  formed,  and  an  acid  reaction  to  test-paper  is  induced  (a). 
But  with  a  trisodium  or  tripotassium  phosphate,  the  solution  remains  neutral 
(6) — «  means  of  distinguishing  the  acid  phosphates  from  the  normal. 

(a)    Na2HP04  +  3AgN03  =  Ag3PO4  +  2NaNO3  +  HNO3 
(6)     Na3PO4  +  3AgNO3  =  Ag3PO4  +  3NaNO3 

Free  orthophosphoric  acid  forms  no  precipitate  with  reagent  silver  nitrate, 
because  silver  phosphate  is  soluble  in  dilute  HNO3  . 

With  lead  acetate  or  nitrate,  Na2HP04  forms  PbaP04 ,  white,  insoluble 
in  acetic  acid,  as  are  also  the  phosphates  of  aluminum  and  ferricum.  With 

*  A  solution  containing  5  p.  c.  ferric  chloride,  mixed  with  one-fourth  its  volume  of  a  10  p.  c. 
soluti'-n  of  orthophosphoric  acid,  requires  that  near  half  of  the  latter  be  neutralized  (so  that 
phosphate  is  to  phosphoric  acid  as  1.114  is  to  1.000)  before  precipitation  occurs.  On  the  other 
hand,  4  cc.  of  a  5  p.  c.  solution  of  ferric  chloride,  mixed  with  1  cc.  of  a  6  p.  c.  solution  of  meta- 
phosphoric  acid,  form  a  precipitate,  to  dissolve  which,  <20  cc.  of  the  same  m;  taphosphoric  acid 
solution  c  r  5  cc.  of  a  24  p.  c.  solution  of  hydrochloric  acid  are  required.  Four  cc.  of  a  5  p.  c. 
solution  of  silvern itrate  with  1  cc.  of  a  10  p.  c.  solution  of  orthophosphoric  acid  give  a  precipi- 
tate, to  dissolve  which  requires  7  cc.  of  the  same  orthophosphoric  acid  solution.  [The  Author's 
report  of  work  by  Mr.  Morgan,  Am.  Jour.  Phar.,  1876,  48,  534.  Kratschmer  and  Sztankovansky, 
ZM  1883,  21,  520.] 


§255,   6.4.  PHOSPHORIC  ACID.  311 

PbCL,  the  precipitate  always  contains  a  chloride.  Free  phosphoric  arid. 
H,P04,  1'onns  an  acid  phosphate,  PbHP04  (Heintz,  Pogg.,  1848,  73,  lli)). 
Lead  salts  also  form  white  precipitates  with  soluble  pyro  and  metaphos- 
phates:  the  pyro  salt,  Pb2P207 ,  is  soluble  in  an  excess  of  Na4P207  .  Bis- 
muth salts  form  BiP04 ,  insoluble  in  dilute  HNO,  . 

Solutions  of  orthophosphates  give,  with  soluble  ferric,  chromic,  and 
aluminum  salts,  mostly  the  normal  phosphates,  FeP04 ,  etc.  The  ferric 
phosphate  is  but  slightly  soluble  in  acetic  acid,  and  for  this  reason  it  is 
made  the  means  of  separating  phosphoric  acid  from  metals  of  the  earths 
and  alkaline  earths  (£152).  Solution  of  sodium  or  potassium  acetate  is 
added;  and  if  the  reaction  is  not  markedly  acid,  it  is  made  so  by  addition 
of  acetic  acid.  Ferric  chloride  (if  not  present)  is  now  added,  drop  by  drop, 
avoiding  an  excess.  The  precipitate,  ferric  phosphate,  is  brownish- white. 

With  zinc  and  manganous  salts,  the  precipitate  is  dimetallic  or  normal— 
ZnHP04 ,  or  Zn3(P04)2 — according  to  the  conditions  of  precipitation. 
When  a  manganic  compound  is  mixed  with  aqueous  phosphoric  acid,  the 
solution  evaporated  to  dryness  and  gently  ignited,  a  violet  or  deep  blue 
mass  is  obtained,  from  which  water  dissolves  a  purple-red  manganic 
hydrogen  phosphate,  a  distinction  from  manganous  compounds.  With  salts 
of  nickel,  a  light  green  normal  phosphate  is  formed;  with  cobalt,  a  reddish 
normal  phosphate. 

Soluble  salts  of  the  alkaline  earth  metals,  with  dimetallic  alkali  phos- 
phates, as  Na2HP04 ,  form  white  precipitates  of  phosphates,  two-thirds 
metallic,  as  CaHP04  ;  with  trimetallic  alkali  phosphates,  white  precipitates 
of  phosphates,  normal  or  full  metallic,  as  Ca.5(P04)2 .  The  precipitates  are 
soluble  in  acetic  acid,  and  in  the  stronger  acids.  Concerning  the  am- 
monium magnesium  phosphate,  see  §189,  6d. 

Magnesium  salts  with  ammonium  hydroxide  give  a  precipitate  of  double 
pyrophosphate,  soluble  in  alkali  pyrophosphate  solution. 

Magnesium  salts  with  ammonium  hydroxide  are  not  precipitated  by 
soluble  metaphosphates  unless  very  concentrated. 

Ammonium  molybdate,  in  its  nitric  acid  solution  (§75,  6d),  furnishes  an 
exceedingly  delicate  test  for  phosphoric  acid,  giving  the  pale  yellow  pre- 
cipitate, termed  ammonium  phosphomolybdate.  The  molybdate  should  be 
in  excess,  therefore  it  is  better  to  add  a  little  of  the  solution  tested  (which 
must  be  neutral  or  acid)  to  the  reagent,  taking  a  half  to  one  cc.  of  the 
latter  in  a  test-tube.  For  the  full  delicacy  of  the  test,  it  should  be  set 
aside,  at  30°  to  40°,  for  several  hours. 
K3P04  +  12(NH4)2MoO4  +  21HNO3  =  (NH4)3PO4.12MoO3  +  21NH,NO3  +  12H2O  . 

Ammonium  molybdate  reacts  but  slowly  with  meta  or  pyrophosphate 
solutions — and  not  until  orthophosphoric  acid  is  formed  by  digestion  with 
the  nitric  acid  of  the  reagent  solution. 


312  PHOSPHORIC  ACID.  §255,    6Z?. 

B. — With  non-metals  and  their  compounds. — Phosphoric  acid  is  not 
reduced  by  any  of  the  reducing  acids.  Phosphates  of  the  first  two  groups 
are  transposed  by  H2S ,  and  of  the  first  four  groups  by  alkali  sulphides 
with  formation  of  a  sulphide  of  the  metal,  except  Al  and  Cr  9  which  form 
a  hydroxide;  phosphoric  acid  or  an  alkali  phosphate  is  also  formed. 
HC1 ,  HN03 ,  and  H2S04  transpose  all  phosphates  and  all  are  transposed 
by  acetic  acid  except  those  of  Pb  ,  Al  and  Fe'"  phosphates.  Sulphurous  acid 
transposes  the  phosphates  of  Ca ,  Mg ,  Mn ,  Ag ,  Pb ,  and  Ba ,  also  the 
arsenite  and  arsenate  of  calcium  (Gerland,  J.  C.,  1872,  25,  39).  Excess  of 
phosphoric  acid  completely  displaces  the  acid  of  all  nitrates,  chlorides,  and 
sulphates  upon  evaporation  and  long-continued  heating  on  the  sand  bath. 

7.  Ignition  with  metallic  magnesium  (or  sodium)  reduces  phosphorus  from 
phosphates  to  magnesium  phosphide,  P2Mg3  ,  recognized  by  odor  of  PH3  , 
formed  on  contact  of  the  phosphide  with  water.  A  bit  of  magnesium  wire  (or 
of  sodium)  is  covered  with  the  previously  ignited  and  powdered  substance  in 
a  glass  tube  of  the  thickness  of  a  straw,  and  heated.  If  any  combination  of 
phosphoric  acid  is  present,  vivid  incandescence  will  occur,  and  a  black  mass 
will  be  left.  The  latter,  crushed  and  wet  with  water,  gives  the  odor  of  phos- 
phorus hydride. 

Orthophosphoric  acid  heated  to  213°  forms  pyrophosphoric  acid;  when  heated 
to  dull  redness  the  meta  acid  is  obtained,  which  sublimes  upon  further  heating 
without  change.  Phosphoric  anhydride,  P2O5  ,  cannot  be  prepared  by  ignition 
of  phosphoric  acid.  Tribasic  orthophosphates,  normal  pyrophosphates,  and 
metaphosphates  of  metals  whose  oxides  are  not  volatile  and  not  decomposed 
by  heat  alone  are  unchanged  upon  ignition.  Bimetallic  orthophosphates, 
M'2HP04  ,  are  changed  to  normal  pyrophosphates  upon  ignition;  also  tribasic 
orthophosphates  when  one-third  of  the  base  is  volatile,  e.  #.,  MgNH(PO4  . 
Mono-metallic  or  primary  orthophosphates,  M'H2P04  ,  become  metaphosphates; 
also  secondary  or  tertiary  orthophosphates  when  only  one  atom  of  hydrogen 
is  displaced  by  a  metal  whose  oxide  is  non-volatile,'  e.  g.,  NaN*H4HP04  . 
Acid  pyrophosphates,  M'2H2P,O7  ,  form  metaphosphates.  When  meta  or  pyro- 
phosphates are  fused  with  an  excess  of  a  non-volatile  oxide,  hydroxide  or 
carbonate  the  tertiary  orthophosphate  is  formed  (Watts',  1894,  IV,  106). 

Phosphates  of  Al ,  Cr  ,  Fe  ,  Cu  ,  Co  ,  Ni  ,  Mn  ,  Grl  and  IT  when  heated  to  a 
white  heat  with  an  alkali  sulphate  form  oxides  of  the  metals  and  an  alkali 
tribasic  orthophosphate;  phosphates  of  Ba ,  Sr ,  Ca ,  Mg1 ,  Zn  and  Cd  form 
double  phosphates,  partial  transposition  taking  place  (Derome,  C.  r.,  1879,  89, 
952;  Grandeau,  A.  Ch.,  1886,  (6),  8,  193). 

8.  Detection. — The  presence  of  orthophosphoric  acid  in  neutral  or  acid 
solutions  is  detected  by  the  use  of  an  excess  of  an  ammonium  molybdate 
solution  (§75,  6d).  With  pyro  and  metaphosphoric  acids  no  reaction  is 
obtained  except  as  they  are  changed  to  the  ortho  acid  by  the  reagents 
used.  Disodium  phosphate,  Na2HP04 ,  after  precipitation  with  silver 
nitrate,  reacts  acid  to  test  papers.  With  trisodium  phosphate  the  solu- 
tion is  neutral  (distinction).  Orthophosphates  are  distinguished  from 
pyro  and  metaphosphates  by  the  color  of  the  precipitate  with  silver  nitrate: 
Ag3P04  is  yellow,  Ag4P,07  and  AgPO,  are  white.  Also  by  the  fact  that 
only  the  ortho  acid  is  precipitated  by  ammonium  molybdate.  Nearly  all 
pyrophosphates  are  soluble  in  sodium  pyrophosphate,  Na4P20T  (("listing- 


£256,  3.  SULPHUR.  313 

lion  from  orthophosphates).  Hager  (J.  C.,  1873,  26,  940)  gives  a  method 
for  detecting  the  presence  of  H3P03 ,  H3As03 ,  or  HNO.{  in  H,P04 .  Sodium 
metaphosphate  does  not  give  a  precipitate  with  ZnS04  cold  and  in  excess; 
with  Na4P20-  and  Na2H,P,07  a  white  precipitate  of  Zn2P207  is  obtained 
(Knorre,  Z.  angew.,  1892,  639). 

9.  Estimation. — (a)  By  precipitation  as  magnesium  ammonium  phosphate, 
MgNH4P04  ,  and  ignition  to  the  pyrophosphate.  (I))  By  precipitation  and 
weighing  as  lead  phosphate,  Pb3(P04)2  .  (c)  By  precipitation  from  neutral  or 
acid  solution  by  ammonium  molybdate  and  after  drying  at  140°  weighing  as 
ammonium  phosphomolybdate.  Consult  Janovsky  (J.  C.,  1873,  26,  91)  for  a 
review  of  all  the  old  methods. 


§256.  Sulphur.     S  =  32.06  .    Usual  valence  two,  four  and  six  (§14). 

1.  Properties. — Sulphur  is  a   solid,   in  yellow,   brittle,   friable  masses    (from 
melt  inn);  or  in  yellowish,  gritty  powder  (from  sublimation)  or  in  nearly  white, 
slightly    cohering,    finely   crystalline    powder    (by   precipitation    from    its    com- 
pounds).    At  — 50°  it  is  white   (Schoenbein,  ./.  pr.,  1852,  55,  101).     The  si>cH/i<- 
gravity  of  native  sulphur  is  2.0748  (Pisati,  B.,  1874,  7,  361).     Melting  point  rhombic 
114.5°.     Boiling   point,    444.53°    (Callendar   and   Griffiths,    C.   N.,    1891,   6  ,    2). 
Vapor  density  at  1160°  is  34,  indicating  that  the  molecule  is  S2  (Bineau,  C.  r.,  1859, 
49,  799);    but  at  lower  temperatures  the  molecule  seems  to  vary  from  £2  to   ;  8. 
Sulphur  is  polymorphous,  existing  in  various  crystalline  forms,  rhombic,   mono- 
clinic  and  triclinic  systems,  and  also  in  amorphous  conditions.     It  is  also  classified 
by  the  relative  solubilities  of  the  various  forms  in  carbon  disulphide.     In  chemical 
activity,  volatility  and  other  properties  it  stands  as  the  second  member  of  the 
Oxygen  Series:    O,  16,000;  S,    32.07;   Se,  79.2;  and  Te,  127.5.     On  being  heated 
it  melts  at  114.5°  to  a  pale  yellow  liquid;    as  the  temperature  rises  it  grows  darker 
and  thicker,  until  at  about  180°  it  is  nearly  solid,  so  that  the  dish  may  be  inverted 
without  spilling.     At  260°  it  again  becomes  a  liquid  as  at  first;    and  at  444.53°  it 
boils  and  is  converted  into  a  brownish-red  vapor.     If  it  is  slowly  cooled,  exactly 
the  same  physical  changes  take  place  in  the  reverse  order,  becoming  thick  at  180° 
and  thin  again  at  114.5°,  and  at  lower  temperatures  solid.     If,  at  a  temperature 
near  its  boiling  point,  it  is  poured  into  cold  water,  it  forms  a  soft,  ductile,  elastic 
string,  resembling  india-rubber.     In  a  few  hours  this  ductile  sulphur  changes  back 
to  the  ordinary  form,  the  change  evolving  heat.      But  if  poured  into  water  from  the 
ether  liquid  form — that  is,  at  114.5° — it  forms  only  ordinary,  brittle  sulphur.     In 
contact  with  air  sulphur  ignites  at  248°  (Hill,  C.  N.,  1890,  61,  125);   burning  in  air 
or  oxygen  with  a  pale  blue  flame  and  penetrating  odor  to  SO2  . 

The  isolated  oxides  of  sulphur  are  S02  ,  S03  ,  S2O8  and  S2O7  .  Sulphur  and 
oxygen  combine  directly  to  form  SO2  and  SO3;  the  former  by  burning  sulphur 
in  oxygen,  the  latter  by  the  action  of  ozone  upon  SO,;  also  by  burning  sulphur 
with  oxygen  under  several  atmospheres  pressure.  S203  is  made  by  dissolving 
sulphur^in  sulphur  dioxide;  S2O7  by  the  action  of  the  electric  discharge  upon 
a  mixture  of  SO3  and  O  . 

2.  Occurrence. — (a)   Found  in  a  free  state,  and  as  SO,  in  volcanic  districts. 
(&)   As  H2S  in  some  mineral  springs,     (c)   As  a  sulphide:   iron   pyrites,   FeS_, ; 
copper   pyrites,    CuFeS2;    orpiment,    As2S3;    realgar,   As.,S2;    zinc    blende,    ZnS; 
cinnabar,  HgS;  galena,  PbS.     (d)   As  a  sulphate:  gypsum,  CaSO4.2ELO;  heavy 
spar,  BaSO4;   kieserite,  MgS04,H2O;  bitter  spar    (Epsom   salts),   MgSO4,7H2O; 
Glauber  salt,  »Ta2S04,10H2O  ,  etc. 

3.  Formation. — (a)  By  decomposing  polysulphides  with  HC1  (Schmidt,  Phar- 
tnnwutiscJic  Chortle,  1898,  175).     (&)  By  adding  an  acid  to  a  solution  of  a  thio- 
sulphate.     (e)   By  the  reaction  between  SO2  and  H2S:  2S02   +   4H2S  =  3S2    -f 
4H2O  .     (d)  By  the  decomposition  of  metallic  sulphides  with  nitric  acid;  2Bi3S, 
+  16HN03  =  4Bi(N03),  +  3S2  +  4NO  +  8H2O  . 


314  suLPHrR.  §256,  4. 

4.  Preparation. —  (a)  The  native  sulphur  is  separated  from  the  clay  and  rock 
in  which  it  is   embedded,  partly   by   melting-   and   partly    by   distillation.     (&) 
From  FeS2  by  heating-  in  close  cylinders  3FeS2  =  Fe3S4  +  S2;  or  at  a  higher 
temperature:  2FeS2  =  2FeS  +  S2  .     Much  of  the  sulphur  contained  in  pyrite* 
is  converted  into  and  utilized  as  sulphuric  acid. 

5.  Solubilities. — Ordinary  (not  precipitated)  sulphur  is  soluble  in  carbon  di- 
sulphide;  the  ductile  variety  is  insoluble.     There  are  several  allotropic  forms 
of  sulphur.     Samples  of  commercial  sulphur  are  almost  never  found  which  are 
entirely  soluble  or  insoluble  in  carbon  disulphide.     Forms  of  sulphur  insoluble 
in    CS2    are    changed    to    soluble    forms    upon    heating    to    the    melting    point; 
also  amorphous  sulphur  insoluble  in  CS2   (formed  by  adding  acids  to  thiosul- 
phates  or  SO2  to  H2S)  is  changed  to  the  soluble  form  by  mixing  with  a  solution 
of  H2S  in  water.     It  dissolves  readily  in  hot  solutions  of  the  hydroxides  of 
potassium,    sodium,    calcium    or   barium,    forming    polysulphides    and   thiosul- 
phates:  3Ca(OH)2  +  6S2  =  2CaS5  +  CaS2O3  +  3H2O  .     These  can  be  separated 
by  alcohol,  in  which  the  sulphides  dissolve.     These  products  are  also  readily 
decomposed  by  acids  with  separation  of  sulphur   (method   of  preparation   of 
precipitated  sulphur). 

Precipitated  sulphur  (in  analysis,  HC1  upon  (NH4)2SX)  is  soluble  in  benzol  or 
low  boiling  petroleum  ether;  of  value  in  analysis  for  the  removal  of  the  sulphur 
to  detect  the  presence  of  traces  of  As  or  Sb  sulphides  (Fresenius,  Z.,  1894,  33, 
573). 

6.  Reactions.  A. — With  metals  and  their  compounds.. — Sulphur  does 
not  combine  with  metals  without  the  aid  of  heat  (see  7),  except  that  under 
very  great  pressure  (6500  atmospheres)  it  combines  with  Pb  ,  Sn  ,  Sb  ,  Bi , 
Cu ,  Cd ,  Fe ,  Zn ,  and  Mg  (Spring,  B.,  1883,  16,  999). 

Flowers  of  sulphur  boiled  with  SnCL  gives  SnS  and  SnCl4  ;  with  HgN03 
almost  exactly  one-half  of  the  mercury  is  precipitated  as  HgS  .  No  action 
with  sulphates  of  Cd ,  Fe",  Mn",  Ni  and  Zn  ;  with  acid  solutions  of  SbCl3 
and  BiCl3 ;  or  with  solutions  of  Asv  and  As'"  (Vortmann  and  Padberg, 
B.,  1889,  22,  2642).  Sulphur  boiled  with  hydroxides  of  K  ,  Na  ,  NH4 ,  Ba , 
Ca,  Sr,  Mg,  Co,  Ni,  Mn,  Hg",  Bi,  Cu',  Cu",  Cd ,  Pb ,  Ag ,  and  Hg' 
forms  sulphides  and  thiosulphates;  also  some  sulphates  are  formed.  No 
action  with  hydroxides  of  Fe,  Zn  and  Sn  (Senderens,  BL,  1891,  (3),  6, 
800). 

B. — With  non-metals  and  their  compounds. 

1.  HCN  warmed  with  sulphur  or  a  polysulphide  becomes  a  thiocyanate: 
2KCN  +  S2  =  2KCNS  or  4HCN  +  2(NH4)2S4  =  4NH4CNS  +  2H2S  +  S2 . 

2.  HN03  becomes  NO  and  H2S04 .     Strong  acid  and  long  continued 
boiling  are  necessary  to  the  complete  oxidation  of  the   sulphur.     The 
crystallized   variety   is   attacked   with   much   greater   difficulty   than   the 
amorphous  or  flowers  (Saint-Gilles,  A.  Ch.,  1858,  (3),  54,  49). 

3.  Red  phosphorus  combines  readily  at  ordinary  temperature,  forming 
P2S3  or  P2S5 ,  depending  upon  the  relative  amounts  of  the  elements  used. 
Ordinary  phosphorus  combines  explosively.    See  §252,  6.    Tribasic  sodium 
or  potassium  phosphate  when  boiled  with  sulphur  forms  alkali  polysul- 
phide and  thiosulphate,  changing  the  phosphate  to  dibasic  phosphate 
(Filhol  and  Senderenj,  (7,  r,,  1883,  96, 1051), 


§257,  1.  HYDROSL'Li'in  me  A<:I/>.  315 

4.  H2S04 ,  concentrated  and  hot,  becomes  S02  from  both  the  S  and  the 
H2S04  :    4H2S04  +  S2  ==  6S02  +  4H20  .     803"  when  added  to  S  at  12° 
forms  the  blue  hyposulpburous  anhydride,  S203   (not  the  anhydride  of 
thiosulphuric  acid,  S202).     S02  reacts  with  S  even  at  ordinary  tempera- 
tures, forming  thiosulphuric  acid  and  tri  or  tetrathionic  acid  (Colefax, 
J.  C.,  1892,  61,  199). 

5.  Cl  in  presence  of  water  forms  HC1  and  H2S04 .     HC103  becomes  HC1 
and  H2S04 . 

6.  Br  in  presence  of  water  becomes  HBr  and  H2S04 .     HBr03  becomes 
HBr  and  H2S04 . 

7.  Sulphur  does  not  appear  to  have  any  action  upon  iodine  or  upon 
iodine  compounds. 

7.  Ignition.— In  the  air,  at  ordinary  temperatures,  finely  divided  sulphur  is 
very  slightly  oxidized,  by  ozone,  to  sulphuric  acid;  at  248°  it  begins  to  oxidize 
rapidly  to  sulphurous  anhydride,  burning  with  a  blue  flame. 

Sulphur,  when  fused  with  the  following  elements,  combines  with  them  to 
form  sulphides:  Pb  ,  Ag  ,  Hg  ,  Sn  ,  As,  Sb  ,  Bi  ,  Cu  ,  Cd  ,  Zn  ,  Co,  Ni  ,  Fe  , 
Sr,  Ca,  Mg,  K,  Na  ,  In,  Tl ,  Pt ,  Pd  ,  Rh  ,  Ir  ,  Li,  Ce  ,  La,  Ne  ,  Pr  . 

Svi— n  becomes  Svi  when  fused  with  alkaline  carbonate  and  nitrate  or  chlorate. 
That  is,  free  sulphur,  S°  ,  or  any  compound  containing  sulphur  with  valence 
less  than  six,  is  oxidized  to  a  sulphate  if  fused  with  an  alkaline  nitrate  or 
chlorate,  nitric  oxide  or  a  chloride  being  formed  and  carbon  dioxide  escaping. 

8.  Detection. — (a)  By  burning  in  the  air  to  a  gas  having  the  odor  of 
burning  matches.     (&)  By   its   solubility   in   CS2 .     (c)  By   formation   of 
H2S04  with  oxidizing  agents,     (d)  By  the  formation  of  sulphides  upon 
fusion  with  metals,     (e)  By  the  blackening  of  silver  coin  after  boiling 
with   alkali  hydroxide,     (f)     Formation   of   reddish-purple  with   sodium 
nitroferricyanlde   after   boiling   with   alkali   hydroxide,      (g)  In   organic 
compounds  by  heating  with  Na  and  testing  the  Na2S  with  sodium  nitro- 
ferricyanide  (Vohl,  B.',  1876,  9,  875). 

9.  Estimation. — Sulphur  is  usually  estimated  by  oxidation  to  a  sul- 
phate and  weighing  as  BaS04 . 


§257.  Hydrosulphuric  acid.     H2S  =  34.076. 
H'2S-",  H  —  S  —  H  . 

1.  Properties. — Molecular  weight,  34.070.  Vapor  density,  17.  Boil  in g  point, 
— 61.8°.  Freezing  point,  — 85.56°.  Under  a  pressure  of  14.G  atmospheres  it  be- 
comes a  liquid  at  11.11°  (Faraday,  A.,  1845,  56,  156).  It  is  a  colorless  poisonous 
gas.  It  burns  readily,  forming  sulphur  dioxide  and  water:  2H2S  +  SO,,  =  2SO2 
+  2H20  .  The  aqueous  solution  slowly  decomposes  upon  exposure  to  the  air 
with  separation  of  sulphur.  The  gas  is  readily  expelled  from  its  aqueous 
solution  by  boiling;  slowly  when  exposed  at  ordinary  temperature.  Both  the 
gas  and  the  water  solutions  have  a  feebly  acid  reaction  towards  moist  litmus 
paper.  They  also  possess  a  strong  characteristic  odor,  resembling  that  of 
rotten  eggs.  In  acid  or  in  alkaline  solutions  it  is  a  strong  reducing  agent. 
See  6. 


HYDROSULPHURIC  ACID.  §257,  2. 

2.  Occurrence. — Found    free    in    volcanic    gases    and    frequently    in    mineral 
springs.     While  the  inhaled  gas  is  poisonous,  the  mineral  waters  containing  it 
are  reputed  to  be  a  healthful  beverage. 

3.  Formation  of  Hydrosulphuric  Acid. — (a)  By  direct  union  of  the  elements 
when  passed  over  pumice  stone  heated  to  400°   (Corenwinder,  A.  Cli^  1852,   (3), 
33:,  77).     (6)  Heating  paraffin  or  tallow  with  sulphur  (Fletcher,  C.  X.,  1879,  40, 
154);  and  by  passing  illuminating  gas  through  boiling  sulphur   (Taylor,  ('.  A., 
1883,  47,   145).     (c)   The  sulphur  in  coal  becomes  H2S  in   the   process  of  gas- 
making,     (d)  From  steam  and  sulphur  at  440°.     (e)  Often  occurs  in  nature  from 
reduction  of  gypsum  by  decaj'ing  organic  matter  (Myers,  J.  pr.,  1869,  108,  123). 
(f)   Transposition  of  sulphides  by  hydracids  or  by  dilute  phosphoric  or  dilute 
sulphuric  acid,     (y)  Decomposition  of  organic  compounds  containing  sulphur. 

Formation  of  Sulphides. — (1)  By  fusion  of  the  metals  with  sulphur,  see 
§256,  7.  (2)  By  action  of  H2S  upon  the  free  metals,  hydrogen  being  evolved. 
With  Hg-  and  Ag  this  occurs  at  ordinary  temperature,  but  with  most  metals  a 
higher  temperature  is  needed.  (3)  Action  of  BLS  on  metallic  oxides  or 
hydroxides.  Those  sulphides  which  are  decomposed  bj-  water  (e.  (/.,  A1,S3  , 
Cr,S3)  are  not  formed  in  its  presence,  but  by  action  of  H2S  upon  the  oxide  at 
a  red  heat.  (4)  By  action  of  soluble  sulphides  upon  metallic  solutions.  The 
ordinary  sulphides  of  the  first  four  groups  are  formed  thus,  except  ferric  salts, 
which  are  precipitated  as  FeS  ,  and  aluminum  and  chromic  salts  as  hydroxides. 
(5)  By  action  of  CS2  upon  oxides  at  a  red  heat.  (6)  By  action  of  free  sulphur 
upon  oxides  at  a  red  heat.  (7)  By  the  action  of  charcoal  upon  the  oxyacids  of 
sulphur  at  a  red  heat  in  presence  of  an  alkaline  carbonate.  To  prepare  a 
sulphide  absolutely  arsenic  free,  take  BaSO4  ,  100  grams;  coal,  pulverized,  25 
grams;  and  NaCl,  20  grams,  mix,  ram  into  a  clay  crucible  and  ignite  to  a 
white  heat  for  several  hours  (Winkler,  Z.,  1888,  27,  26).  (8)  By  the  action  of 
zinc  amalgam  on  sulphuric  acid  (Walz,  C.  A'.,  1871  23,  245).  (9)  As  a  reagent 
for  the  formation  of  metallic  sulphides  in  analysis  it  is  recommended  by 
Schiff  and  Tarugi  (#.,  1894,  27,  3437),  Schiff  (£.,  1895,  28,  1204),  and  Tarugi 
(Gazzetta,  1895,  25,  i,  269),  to  use  ammonium  thioacetate,  CH3COSNH4;  prepared 
by  distilling  a  mixture  of  phosphorus  pentasulphide  and  glacial  acetic  a«id 
(300  grams  each)  with  150  grains  of  cracked  glass.  A  large  distilling  flask  is 
used  and  the  distillate  is  collected  to  103°.  It  is  then  dissolved  in  a  slight 
excess  of  ammonium  hydroxide,  diluting  to  three  volumes  from  one  volume 
of  the  acid.  Salts  of  the  metals  of  the  first  two  groups  in  acid  solution  are 
readily  precipitated  as  sulphides  upon  warming  with  this  reagent. 

1.  2Fe  +  S2  =  2FeS 

2.  2Ag  +  H2S  =  Ag2S  +  H2 

5.    Pb(OH)2  +  H2S  =  PbS  +  2HCO 

4Fe(OH)3  +  6H2S  =  4FeS  +  S2  +  12H20 

4.  4FeCl3  +  6(NH4)2S  =  4FeS  +  S2  +  12NH4C1 

5.  2CaO  +  CS2  —  2CaS  +  CO2 

6.  4CaO  +  3S2  —  4CaS  +  2S02 

7.  K2S04  +  20  =  K2S  +  2C02 

4.  Preparation. — For  laboratory  purposes  it  is  nearly  always  made  by 
adding  H2S04  or  HC1  to  FeS  .     The  ferrous  sulphide  is  prepared  either 
by  fusion  of  the  iron  with  the  sulphur,  or  by  bringing  red  hot  iron  rods 
in  contact  with  sticks  of  sulphur,,  and  is  made  to  drop  into  tubs  of  cold 
water.     Dilute  H,S04  should  be  used:*  FeS  +  H,S04  ==  FeS04   +   HJS . 
Concentrated  H2S04  has  no  action  on  FeS ,  unless  heated  and  then  S02  is 
evolved:   2FeS  +  10H2S04  =  Fe2(S04)3  +  9S02  +  10H20  ;  and  frequently 
free  sulphur  is  formed  by  the  action  of  the  H2S  upon  the  S02  first  formed. 

*  If  the  acid  is  diluted  with  eleven  volumes  of  water  ferrous  sulphate  crystals  will  not  be 
deposited. 


§257,  5.  HYDROSULPHURIC   ACID.  317 

The  colorless  ammonium  sulphide,  (NH4)2S,  is  prepared  by  saturating 
ammonium  hydroxide  with  H2S  until  a  sample  will  no  longer  give  a  pre- 
cipitate with  a  solution  of  magnesium  sulphate;  showing  that  ammonium 
hydroxide  is  no  longer  present.  Upon  standing  the  solution  gradually 
becomes  yellow  with  formation  of  the  poly  sulphides  or  yellow  ammonium 
sulphide,  (NH4)2SX  This  may  be  hastened  by  the  addition  of  sulphur 
(Bloxam,  J.  C.,  1895,  67,  277). 

Sodium  sulphide,  Na2S ,  is  prepared  by  neutralizing  an  alcoholic  solution 
of  NaOH  with  H2S  and  then  adding  an  equal  amount  of  NaOH  and  allowing 
to  crystallize ;  air  being  excluded.  The  various  polysulphides,  Na2S2  to 
Na2S5 ,  arc  prepared  by  boiling  the  normal  sulphide  with  the  calculated 
amounts  of  sulphur  (Boettger,  A.,  1884,  223,  335;  Geuther,  A.,  1884,  224, 
201). 

5.  Solubilities. — At  15°  water  dissolves  2.66  volumes  of  the  gas  H2S  . 

Sulphides  which  dissolve  in  dilute  H2S04  evolve  H2S ,  e.  g.,  CdS ,  FeS , 
MnS ,  ZnS ,  etc.  But  if  a  sulphide  requires  concentrated  H2S04  for  its 
solution;  S  and  S02  are  formed  or  S02  alone;  e.  g.,  Bi2S3 ,  CuS,  HgS .  If 
concentrated  H2S04  be  used  upon  a  sulphide  that  might  have  been  dis- 
solved in  the  dilute  acid,  then  no  H2S  is  evolved:  ZnS  +  4H2S04  =±  ZnS04 
+  4S02  +  4H,0  .  Or  with  a  small  amount  of  water  present:  2ZnS  + 
4H2S04  =  2ZnS04  +  S2  +  2S02  +  4H20  .  The  sulphur  of  the  zinc  sul- 
phide is  oxidized  to  free  sulphur  and  that  of  the  sulphuric  acid  is  reduced 
to  sulphur  dioxide.  HgS  is  almost  insoluble  in  HN03 ,  dilute  or  concen- 
trated, readily  soluble  in  chlorine,  nitrohydrochloric  acid,  or  chloric  acid 
if  kot.  Most  other  sulphides  are  soluble  in  hot  HN03  (§74,  6e).  Long 
continued  boiling  with  water  more  or  less  completely  decomposes  the  sul- 
phides of  Ag ,  As ,  Sb ,  Sn ,  Fe ,  Co  ,  Ni ,  and  Mn  ;  no  effect  with  sulphides 
of  Hg,  Au,  Pt,  Mo,  Cu,  Cd,  and  Zn  (Clermont  and  Frommel,  A.  Ch.y 
1879,  (5),  18,  203). 

As  a  reagent,  hydrosulphuric  acid,  gas  or  solution  in  water  finds  ex- 
tended application  in  the  analytical  laboratory.  The  grouping  of  the 
bases  for  analysis  depends  very  largely  upon  the  relative  solubilities  of  the 
sulphides.  Hydrosulphuric  acid  in  alkaline  solution,  alkali  sulphide  or 
polysulphide,  is  a  scarcely  less  important  reagent.,  being  especially  valuable 
in  the  subdivision  of  the  metals  of  the  second  group. 

The  sulphides  of  the  first  four  groups  are  insoluble.  Hydrosulphuric 
acid  transposes  salts  of  the  first  two  groups  in  acid,  neutral,  and  alkaline 
mixtures,  except  arsenic,  which  is  generally  imperfectly  precipitated  un- 
less some  free  acid  or  salt  that  is  not  alkaline  to  litmus  be  present.  The 
result  is  a  sulphide,  but  mercurosum  forms  mercuric  sulphide  and  mer- 
cury, and  arsenic  acid  may  form  arsenous  sulphide  and  free  sulphur, 
Ferric  solutions  are  reduced  to  ferrous  with  liberation  of  sulphur.  In  acid 
mixture  other  third  and  fourth  group  salts  are  not  disturbed,  but  from 


318  HYDROSULPHURIC  ACID.  §257,  6. 

solutions  of  their  normal  salts  traces  of  cobalt,  nickel,  manganese,  and 
zinc  (§135,  60)  are  precipitated. 

Soluble  sulphides  transpose  salts  of  the  first  four  *  groups.  The  result 
is  a  sulphide,  except  that  with  aluminum  and  chromium  salts  it  is  a 
hydroxide,  hydrosulphuric  acid  being  evolved.  With  mercurous  salts, 
mercuric  sulphide  and  mercury  are  formed;  with  ferric  salts,  ferrous  sul- 
phide and  sulphur. 

The  precipitates  have  strongly  marked  colors — that  of  zinc  being  white; 
manganese,  flesh  colored;  those  of  iron,  copper,  and  lead,  black;  arsenic 
stannic  and  cadmium,  yellow;  antimony,  orange-red;  stannous,  brown;  mer- 
cury, successively  white,  yellow,  orange,  and  black. 

6.  Reactions.  A. — With  metals  and  their  compounds,— Some  metals 
are  converted  into  sulphides  on  being  treated  with  hydrosulphuric  acid; 
e.  g.,  Ag ,  Cu ,  Hg ,  etc.  The  alkali  polysulphides  slowly  attack  many 
metals  with  formation  of  sulphides:  Sn  becomes  M'2SnS3  ;  Ag  becomes 
Ag2S,  no  action  with  colorlesc  (NH4)2S ;  Ni  forms  NiS ;  Fe ,  FeS;  Cu, 
CuS  and  then  Cu2S  (with  colorless  ammonium  sulphide,  (NH4)2S ,  Cu2S 
is  formed  with  evolution  of  hydrogen)  (Priwozink,  A.,  1872,  164,  46). 

The  hydroxides  or  non-ignited  oxides  of  Pb",  Ag ,  Hg",  Sb ,  Sn ,  Bi'", 
Cu,  Cd,  Fe",  Co",  Ni",  Mn",  Zn,  Ba,  Sr ,  Ca,  Mg,  K,  Na,  and  NH4 
unite  with  moist  H2S  at  ordinary  temperature  to  form  sulphides  without 
change  of  the  valence  of  the  metal.  In  other  cases  the  valence  of  the 
metal  is  changed,  usually  with  liberation  of  sulphur. 

1.  Pb"+n  becomes  PbS  and  S  . 

2.  Asv  in  acid  solution  forms  some  As2S3  and  S .     See  §69,  6e. 

3.  Hg'  becomes  HgS  and  Hg  . 

4.  Crvl  becomes  Cr'"  and  S ,  if  the  H2S  be  in  excess :   2K,Cr207  +  8H2S 
=  4Cr(OH)3  -f  3S2  +  2K2S  +  2H,0  . 

5.  Fe"'  becomes  Fe"  and  S  :   4FeCl3  +  2H2S  =:  4FeCl2  +  4HC1  +  S2 . 
If  the  solution  be  alkaline  FeS  is  precipitated :  4FeCl3  -f  6K..S  =  4FeS  + 
12KC1  +  S2 . 

6.  Co"+n  becomes  Co"  and  S  . 

7.  Ni"+n  becomes  Ni"  and  S . 

8.  Mn"+n  becomes  Mn"  and  S.     In  alkaline  solution  with  excess  of 
KMn04 ,  an  alkali  sulphate  is  formed  and  Mn02  :    8KMn04  -f  3K2S  = 
3K2S04  +  4K20  +  8Mn02  (Schlagdenhafen,  Bl.,  1874,  (2),  22,  16). 

In  the  above  reactions,  if  an  alkaline  sulphide  be  used  instead  of  hydro- 
sulphuric  acid,  the  metal  will  be  precipitated  as  a  sulphide  with  the 

*  The  normal  fixed  alkali  sulphides  (Na2S,  K2S),  precipitate  solutions  of  calcium  and  mag- 
nesium salts  as  the  hydroxides  :  Ca(C2H3O2^  +  2Na2S  +  2BT2O  =  Ca(OH)3  -f  2XaC2H3O3  + 
2NaHS.  No  reaction  with  the  acid  fixed  alkali  sulphides  (NaHS,  KHS)  or  with  ammonium 
sulphides  (Pelouze,  A.  C/iM  1866,  (4),  7, 172). 


§257,  7.  HYDROSULPHURIC  ACID.  319 

formation  of  an  alkali  hydroxide;  except  that  the  arsenic  will  remain  in 
solution  (§69,  5c)  and  the  chromium  will  be  precipitated  as  the  hydroxide. 
Dry  H2S  has  no  action  on  the  dry  salts  of  Pb ,  Ag ,  Hg ,  As ,  Sb ,  Sn , 
Bi ,  Cu ,  Cd  ,  or  Co  ;  nor  does  it  redden  dry  blue  litmus  (Hughes,  Phil. 
Mag.,  1892,  (5),  33,  471). 

Many  insoluble  sulphides,  freshly  precipitated,  transpose  the  solutions  of 
other  metallic  salts.  In  some  cases  the  action  is  quite  rapid  at  ordinary  tem- 
perature, in  others  long-continued  heating1  (several  hours)  at  100°  is  necessary. 
PdS  is  formed  by  action  of  PdCL  with  sulphides  of  all  the  metals  following  in 
the  series  below  named,  but  PdS  is  not  transposed  by  solutions  of  the  metals 
following-.  Silver  salts  form  Ag2S  with  sulphides  of  the  metals  following  in  the 
series  but  not  with  sulphides  of  Pd  and  Hg ,  etc.;  Pd  ,  Hg  ,  Ag  ,  Cu  ,  Bi  ,  Cd  , 
Sb  ,  Sn  ,  Pb  ,  Zn  ,  Ni  ,  Co  ,  Fe  ,  As  ,  Tl  and  Mn  (Schiirmann,  A.,  1888,  249,  326). 

B. — With  non-metals  and  their  compounds. 

1.  H3Fe(CN)6  becomes  H4Fe(CN)6  and  S.     Proof:    Boil  to  expel  the 
excess  of  hydrosulphuric  acid,  then  add  ferric  chloride  (§126,'  66). 

2.  HN03  becomes  NO  and  S  .     If  the  HN03  be  hot  and  concentrated  the 
sulphur  is  oxidized  to  sulphuric  acid. 

3.  H2S  has  no  reducing  action  on  the  acids  of  phosphorus. 

4-  H2S03  becomes  pentathionic  acid,  H2S506 ,  and  sulphur:  10H2S03  -j- 
10H,S  ==  2H2S506  +  5S2  +  18H20  .  With  excess  of  H2S  the  product  is 
entirely  free  sulphur  from  both  compounds:  2H2S03  +  4H2S  =  3S2  + 
6H20  (Debus,  J.  (7.,  1888,  53,  282). 

H2S04 ,  dilute  no  action;  concentrated  and  hot,  S  and  S02  are  formed: 
2H2S04  +  2H2S  =  S2  +  2S02  +  4H20  ($256,  6J54). 

5.  Cl  with  H2S  in  excess  forms  HC1  and  S  ;  with  Cl  in  excess  forms  HC1 
and  H,S04 . 

HC103  with  H2S  in  excess  forms  HC1  and  S  ;  with  HC103  in  excess  HC1 
and  H2S04 . 

6.  Br  with  H..S  in  excess  forms  HBr  and  S  ;  with  Br  in  excess  HBr  and 
H2S04 . 

HBrO,  with  H2S  in  excess  forms  HBr  and  S  ;  with  HBr03  in  excess  HBr 
and  H,S04 . 

7.  I  becomes  HI  and  S  (Filhol  and  Mellies,  A.  Ch.,  1871,  (4),  22,  58). 
HIO;!  becomes  HI  and  S  . 

7.  Ignition. — Dry  hydrosulphuric  acid  gas  is  not  decomposed  when  heated  to 
350°  to  360°.  At  this  temperature  AsH3  in  presence  of  potassium  polysulphide, 
K2S3  ,  liver  of  sulphur,  is  decomposed:  2AsH3  +  3K2S3  =  2K,AsS3  +  3H2S; 
thus  furnishing  a  ready  means  of  purifying-  H2S  for  lexicological  work  (§69, 
6'6)  (Pfordten,  B.,  1884,  17,  2897). 

If  air  be  excluded  some  sulphides  may  be  sublimed  unchanged;  c.  f/.,  HgS  , 
As.,S8  ,  As2S6  ,  Sb,S8  ,  etc.  In  some  cases  part  of  the  sulphur  is  separated, 
leaving  a  sulphide  of  a  lower  metallic  valence:  2FeS2  =  2FeS  +  S2  .  Some 
sulphides  remain  unchanged  upon  ignition  in  absence  of  air;  e.  g.,  FeS  ,  MnS  , 
CdS  ,  etc.  All  sulphides  suffer  some  change  on  being  ignited  in  the  air;  some 
slowly,  others  rapidly;  Sb2S8  ,  CuS  ,  A12S3  ,  Cr,S3  ,  etc.,  evolve  SO,  and  leave 


£20  HYDROSULPHURIC  AC'lh.  £257,  8. 

the  oxide  of  the  metal;  HgS  ,  Ag-,8  ,  etc.,  evolve  S02  and  leave  the  free  metal. 
All  sulphides,  as  well  as  all  other  compounds  of  sulphur,  when  fused  with  KN03 
or  KC103  in  presence  of  an  alkali  carbonate  are  oxidized  to  an  alkali  sulphate; 
forming-  NO  or  KC1  and  evolving  CO,  .  The  metal  is  changed  to  the  carbonate, 
oxide  or  the  free  metal  (§228,  7). 

When  ignited  on  charcoal  with  sodium  carbonate — or  (distinction  from 
sulphates)  if  ignited  in  a  porcelain  crucible  with  sodium  carbonate — soluble  sodium 
••sulphides  are  obtained.  The  production  of  the  sodium  sulphide  is  proved  by  the 
black  stain  of  Ag2S  ,  formed  on  metallic  silver  by  a  moistened  portion  of  the 
fused  mass.  (Compounds  of  selenium  and  tellurium,  §§112  and  113.) 

8.  Detection.— (a)  The  odor  of  the  gas  constitutes  a  delicate  and  char- 
acteristic test  when  not  mixed  with  other  gases  having  a  strong  odor. 
(&)  The  gas  blackens  filter  paper  moistened  with  a  solution  of  lead  ace- 
tate, delicate  and  characteristic.  In  the  detection  of  traces  of  the  gas, 
a  slip  of  bibulous  paper,  so  moistened,  may  be  inserted  into  a  slit  in  the 
smaller  end  of  a  cork,  which  is  fitted  to  the  test-tube,  wherein  the  material 
to  be  tested  is  treated  with  sulphuric  acid;  the  tube  being  set  aside  in  a 
warm  place  for  several  hours.  If  any  oxidizing  agents  are  present — as 
chromates,  ferric  salts,  manganic  salts,  chlorates,  etc. — hydrosulphuric 
acid  is  not  generated,  but  instead  sulphur  is  separated,  or  sulphates  are 
formed  (6).  (c)  The  gas  blackens  silver  nitrate  solution,  delicate  but 
PH.{ ,  AsH3 ,  and  SbH3  also  blacken  silver  nitrate  solution,  (d)  By  its 
reducing  action  upon  nearly  all  oxidizing  agents  with  separation  of  sul- 
phur, which  is  detected  according  to  §256,  8.  KMn04  is  perhaps  the  most 
delicate  test  but  the  least  characteristic,  (e)  Its  oxidation  to  a  sulphate 
is  characteristic  in  absence  of  other  sulphur  compounds.  This  method 
is  usually  employed  with  sulphides  not  transposed  by  dilute  H2S04  ; 
chlorine,  nitrohydrochloric  acid  or  bromine  being  the  usual  oxidizing 
agents.  Also,  these  sulphides  and  certain  supersulphides,  attacked  with 
difficulty  by  acids,  as  iron  pyrites  and  copper  pyrites,  are  reduced  and 
dissolved,  with  evolution  of  Jiydrosulphuric  add,  by  dilute  sulphuric  acid 
with  zinc.  The  gas,  with  its  excess  of  hydrogen,  may  be  tested  by  method 
(/).  (/)  Sodium  nitroferricyanide  Na2[Fe(CN)5(NO)].2H20  also  known 
as  sodium  nitroprusside  gives  a  very  delicate  and  characteristic  test  for 
H2S  as  an  alkali  sulphide.  The  gas  is  passed  into  ammonium  hydroxide; 
and  to  this  mixture  a  20  per  cent  solution  of  the  reagent  is  added,  pro- 
ducing a  transient  reddish-purple  color.  Free  H2S ,  dilute,  remains 
colorless;  a  concentrated  solution  gives  a  blue  color,  due  to  the  reducing 
action  of  the  H2S  on  the  ferricyanide.  Caustic  alkali  hinders  the  reac- 
tion. 0.000018  grain  of  H2S  ,  as  gas  or  alkali  sulphide,  can  be  detected  by 
this  reagent  (Reichard,  Z.,  43,  222).  (g)  The  most  delicate  test  for  hy- 
drogen sulphide  involves  the  formation  of  methylene  blue  (E.  Fisher,  Ber., 
16,  2234).  By  this  test  0.02  mg.  of  hydrogen  sulphide  in  a  liter  will  give 
a  blue  color  after  standing  half  an  hour.  The  test  is  carried  out  as 
follows :  to  the  solution  to  be  tested  is  added  one-tenth  of  its  volume 


§258,  5.  THIOSULPHURIC    ACID.  321 

of  concentrated  HC1,  then  a  small  amount  of  dimethylparaphenylendiamine 
sulphate  (NH2.C0H4.N(CH3)2.H2S04)  and  after  it' has  dissolved  a  drop 
or  two  of  dilute  ferric  chloride  solution. 

For  method  of  separation  of  the  various  sulphur  compounds  from  each 
other  consult  Kynaston  (J.  C.,  1859,  11,  166),  Bloxam  (C.  N.,  1895, 
72,  63),  Votocek  (Ber.,  1907,  40,  414)  and  Autenrieth  and  Windaus 
(Z.,  1898,  295). 

9.  Estimation.— Sulphides  are  usually  oxidized  to  H2S04  (by  chlorine, 
bromine,  or  nitrohydrochloric  acid,  or  by  fusion  with  KN03  and  Na2C03) 
precipitated  with  BaCl2  and  weighed  as  BaS04 . 


§258.  Thiosulphuric  acid.     H2S203  =  114.136. 

Difhionous  acid. 

0 

II 
H'2(S2)IV0-"3 ,  H  —  0  —  S  —  S  —  H.* 


1.  Properties.— Thiosulphuric  acid,  H2S2O3    (formerly  called  hyposulphurous 
acid),  has  not  been  isolated;  but  it  almost  certainly  exists  in  dilute  solutions, 
when  a  dilute  weak  acid  is  added  to  a  solution  of  sodium  thiosulphate,  Na2S203  , 
soon  beginning  to  decompose  into  H2S03   and  S   (Landolt,  B.,  1883,   16,  2985). 
The  thiosulphates  are  not  particularly  stable  compounds,   some  decomposing 
almost  immediately  upon  forming;   e.  g.,  mercury   thiosulphates.     Alkali   thio- 
sulphates decompose  upon  heating  into  sulphate  arid  polysulphide:  4Na2S2O3  == 
3Na,S04   +  NsLjSs  .     Other  salts  give  also  S  and   H2S  .     Boiling  solution  of  a 
thiosulphate  gives  a  sulphate  and  H2S  or  a  sulphide  of  the  metal. 

2.  Occurrence. — Not  found  in  nature. 

3.  Formation. — Thiosulphates    are    formed    by    the    oxidation    of    alkali    or 
alkaline  earth   polysulphides   by   exposure   to  the   air   or  by   SO2    or   K2Cr,,07: 
2CaS5  +  302  =  2CaS,03  +  3S2;  4Na2S5   +  6S02  =  4Na,S2O3   +  9S2;  2K,S5   + 
4K2Cr207   +  13H2O  =  5K2S2O3   +  8Cr(OH)3   +  2KOH   (Doepping,  A.,  1843,  46, 
172;  Gueront,  C.  r.,  1872,  75,  1276).     Also  by  heating  ammonium  sulphate  with 
phosphorus  pentasulphide  (Spring,  /?.,  1874,  7,  1157). 

4.  Preparation. — Thiosulphates  are  prepared  by  boiling   sulphur  in  a   solu- 
tion of  normal  alkali   sulphite:  2Na2SO3    +   S2  =  2Na,S2O3  .     Fixed   alkali   or 
alkaline  earth  hydroxides  with  sulphur  also  form  thiosulphates:  3Ca(OH)o   + 
6S2   =  2CaS5    +  'CaS2O3    +    8H,O    (Filhol   and   Senderens,   C.   r.,    1883,   96,   839; 
Senderens,  C.  ?\,  1887,   104,  58).     Commercial  sodium  thiosulphate  is  prepared 
by  passing  SO,  into  "  soda  waste  "  suspended  in  water,  calcium  thiosulphate 
being  formed.     This  is  treated  with  sodium  sulphate,  filtered  and  evaporated 
to  crystallization. 

5.  Solubilities. — The  larger  number  of  the  thiosulphates  are  soluble  in  water; 
those  of  barium,  lead  and  silver  being  only  very  sparingly  soluble.     The  thio- 
sulphates are  insoluble  in  alcohol.     They  are  decomposed,  but  not  fully  dis- 
solved, \>y  acids,  the  decomposition  leaving  a  residue  of  sulphur. 

*  Bunte,  B.,  1874,  7,  646. 


T'HIOSVLPH'URIC  ACID.  §258,  (5. 

Alkali  thiosulphate  solutions  dissolve  the  thiosulphates  of  lead  and  silver; 
also  the  chloride,  bromide  and  iodide  of  silver,  and  mercurous  chloride;  th.6 
iodide  and  sulphate  of  lead;  the  sulphate  of  calcium,  and  some  other  precipi* 
tates — by  formation  of  soluble  double  thiosulpliatcs: 

Ag2S203  +  Na2S203  =  2NaAgS20. 
AgCl  +  Na2S203  =  NaAgS203  +  NaCl 
PbS04  +  3Na2S203  =Na4Pb(S203)3  +  Na2S04 

6.  Reactions. — A. — "With  metals  and  their  compounds. — With   soluble  thio- 
sulphates, solutions  of  lead  and  silver  salts  are  precipitated  as  thiosulphates, 
white,  soluble  in  excess  of  alkali  thiosulphate.     These  precipitates  decompose 
upon  standing,  rapidly  on  warming,  into  sulphides  and  sulphuric  acid:  Ag.S.Oj 
-f-  H,O  =  AgoS  -f  H.,S04  .     Soluble  mercury  salts  with  sodium  thiosulphate 
fornTa  white  precipitate,  almost  instantly  turning  black  with  decomposition  to 
mercuric  sulphide.     Na2S,O3  blackens  HgCl  ,  a  portion  of  the  mercury  going 
into  solution,  colorless,  reprecipitated  black  upon  warming. 

Acid  solutions  of  arsenic  and  antimony  are  precipitated  by  hot  solution  of 
Na2S2O3  as  sulphides,  As2S3  and  Sb2S3  (a  separation  from  tin,*  which  is  not 
precipitated)  (6e,  §§69,  70  and  71).  Solutions  of  copper  salts  with  thiosul- 
phates, on  long  standing,  precipitate  cuprous  salt,  changed  by  boiling  to 
cuprous  sulphide  and  sulphuric  acid  (separation  from  cadmium,  §78,  Ge). 

Solutions  of  ferric  salts  are  reduced  to  ferrous  salts  with  formation  of  sodium 
tetrathionate:  2FeCl3  +  2Na2S2O3  =  SFeCL  +  2NaCl  +  Na2S4O8;  used  as  a 
quantitative  method  of  estimation,  with  a  few  drops  of  potassium  thiocyanate 
as  an  indicator.  Chromic  acid  (chromates  in  acid  solution)  are  reduced  to 
chromic  salts  with  oxidation  of  the  thiosulphate. 

Permanganates  in  neutral  solution  become  manganese  dioxide,  in  acid  solu- 
tion the  reduction  is  complete  to  manganous  salt,  a  sulphate  and  dithionate 
being  formed  (Luckow,  Z.,  1893,  32,  53). 

Barium  chloride  forms  a  white  precipitate  of  barium  thiosulphate,  BaS203  , 
nearly  insoluble  in  water;  100  parts  of  water  dissolve  0.2675  part  of  BaS2O3  H2O 
at  17.5°.  Calcium  chloride  forms  no  precipitate  (distinction  from  a  sulphite). 

B. — With  non-metals  and  their  compounds. — When  thiosulphates  are  decom- 
posed by  acids,  the  constituents  of  thiosulphuric  acid  are  dissociated  as  sul- 
phurous acid  and  sulphur.  Nearly  all  acids  in  this  way  decompose  thiosul- 
phates: 2Na2S2O3  +  4HC1  =  4NaCl  +  2H2S03  -f  S,  . 

Thiosulphates  are  reducing  agents — even  stronger  and  more  active  than  the 
sulphites  to  which  they  are  so  easily  converted.  This  reduction  is  illustrated 
by  the  action  on  arsenic  compounds,  on  ferric  salts  and  on  chromates  and 
permanganates  as  given  above.  Also  the  halogens  are  reduced  to  the  halide 
salts  forming  a  tetrathionate:  2Na2S203  +  I2  =  2NaI  +  Na2S4O6  .  If  chlorine 
or  bromine  be  in  excess  the  tetrathionate  is  further  oxidized  to  a  sulphate: 
Na2S203  +  4CL  -f  5H20  =  Na2S04  +  H2S04  +  8HC1 .  Chloric,  bromic  and 
iodic  acids  are  first  reduced  to  the  corresponding  halogens  and  then  with  an 
excess  of  the  thiosulphate  to  the  halides,  always  accompanied  with  the  separa- 
tion of  sulphur.  Nitric  acid  is  reduced  to  nitric  oxide  wTith  the  separation  of 
sulphur. 

7.  Ignition.— On  ignition,  or  by  heat  short  of  ignition,  all  thiosulphates  are 
decomposed.     Those  of  the  alkali  metals  leave  sulphates  and  poly  sulphides  (a), 
others    yield    sulphurous    acid    with    sulphides,    or    sulphates,  "or    both.     The 
capacity    of    thiosulphates    for    rapid    oxidation,    renders    their    mixture    with 
chlorates,  nitrates,  etc.,  explosive,  in  the  dry  way.     Chlorates  with  thiosulphates 
explode  violently  in  the  mortar.     Cyanides  and  ferricyanides,  fused  with  thin- 
sulphates,  form  thiocyanates,  which  may  be  dissolved  by  alcohol  from  other 
products.     By  fusion  on  charcoal  with  Na,CO3  ,  thiosulphates  form  sulphides 
(&)   and   (c);  and  by  fusion  with  an  alkali  carbonate  and  nitrate  or  chlorate, 

*  According  to  Vortmann  (M.,  1886,  7,  418)  sodium  thiosulphate  may  be  used  instead  of  hydro- 
sulphuric  acid  in  the  second  group  of  bases.  An  excess  of  the  reagent  is  to  be  avoided  and 
nitric  acid  should  be  absent. 


§259.  HYPOSULPHUROUS   ACID.  323 


a  sulphate  is  formed  (d).  By  ignition  of  a  metallic  salt  with  Na^Os  in  a 
dry  test-tube  the  characteristic  colored  sulphide  of  the  metal  is  obtained 
(Landauer,  B.,  1872,  5,  406). 

(a)  4Na2S203  =  Na2S5  +  3Na2S04 

(&)       Na2S203  +  ETa2C03  +  2C  =  2Na2S  +  3C02 

(c)  2PbS203  +  4Na2C03  +  5C  =  4Na2S  +  2Pb  +  9C02 

(d)  3Na2S203  +  3Na,C03  +  4KC103  =  6Na2S04  -f  4KC1  +  3C02 

8.  Detection.  —  In  analysis,  thiosulphates  are  distinguished  by  giving  a  pre- 
cipitate of  sulphur  with  evolution  of  sulphurous  anhydride  when  their  solu- 
tions are   treated   with   hydrochloric   acid;   by   their   intense    reducing   power, 
shown  in  the  blackening  of  the  silver  precipitate;  and  by  non-precipitation  of 
calcium  salts. 

The  precipitation  of  sulphur  with  evolution  of  sulphurous  anhydride,  by  addition 
of  dilute  acids  —  as  hydrochloric  or  acetic  —  is  characteristic  of  thiosulphates. 
It  will  be  understood,  however,  that  in  presence  of  oxidizing  agents,  which  can 
be  brought  into  action  by  the  acid,  sulphides  wrill  likewise  give  a  precipitate  of 
sulphur. 

In  the  presence  of  a  sulphate  and  sulphite  the  thiosulphate  is  detected  as 
follows:  Add  BaCl2  and  NH4C1  in  excess,  then  HC1  to  solution  of  all  but  the 
BaSO.i  .  Filter  and  treat  the  filtrate  with  iodine,  forming  BaSO4  of  the  sulphite 
and  BaSiO,,  of  the  thiosulphate.  Filter  and  add  bromine  to  the  filtrate,  which 
then  forms  BaS04  (Smith,  C.  N.,  1895,  72,  39). 

Sulphides,  sulphites  and  thiosulphates  may  be  separated  as  follows:  —  Add  to  the 
neutral  solution  cadmium  carbonate,  shake  and  filter  off  the  cadmium  sulphide  and 
excess  of  cadmium  carbonate.  Test  the  filtrate  for  sulphide  with  sodium  nitro- 
prusside  and  again  add  cadmium  carbonate  until  the  sulphide  is  entirely  removed. 
Add  strontium  nitrate  to  the  filtrate,  allow  to  stand  overnight  and  filter  off  the 
strontium  sulphite.  Test  the  precipitate  for  sulphurous  acid.  Test  the  filtrate 
for  thiosulphuric  acid  by  acidifying  with  HC1  and  warming.  (Autenrieth  and 
Windaus,  Z.,  1898,  295.) 

9.  Estimation.  —  By  titration  with  a  standard  solution  of  iodine,  or  by  titrating 
the  iodine  liberated  by  a  standard  solution  of  potassium  dichromate  (§§125,  10, 
and  279,  657). 

§259.  Hyposulphurous  acid,    H2S02  =  66.076. 


(Hydro sulphurous  or  diihionous  acid.) 

H'2S"-0-"2,  H  — 0  — S  — H. 

II 
0 

Obtained  by  Schiitzenberger  (G.  r.,  1869,  69,  196)  by  the  action  of  zinc  on 
sulphurous  acid:  Zn  +  2SO2  +  H2O  —  ZnSO  +  H2SO2  .  The  sodium  salt  is 
formed  by  treating  a  concentrated  solution  of  sodium  acid  sulphite  with  zinc 
filings:  Zn  +  3NaHSO3  =  ZnSO3  +  Na2SO3  +  NaHSO2  +  H2O  .  In  the  forma- 
tion of  the  free  acid  or  of  the  sodium  salt  no  hydrogen  is  evolved.  It  is  a  very 
unstable  compound,  a  strong  reducing  agent,  rapidly  absorbs  oxygen  from  the 
air,  becoming  sulphurous  acid  or  a  sulphite.  According  to  Bernthsen  (B.,  1881, 
14,  438)  the  sodium  salt  does  not  contain  hvdrogen.  He  gives  the  formula  as 
Na2S204:  Zn  +  4NaHSO3  =  ZnS03  +  Na2SO3  +  Na2S2O4  +  2H2O  .  It  is  used 
in  the  preparing  of  indigo  white  for  the  printing  of  cotton  fabrics.  See  also 
Dupre,  J.  C.,  1867,  20,  291, 


324  DITH IONIC    ACID—TRITHIONIC   ACID.  §260. 

§260.  Dithionic  acid.     H2S,06  =  162.136. 

0       0 

II       II 

H'2(S2)xO-"6 ,  H  —  0  —  S  —  S  —  0  —  H  . 
II       II 
0       0 

Known  only  in  the  form  of  its  salts  and  as  a  solution  of  the  acid  in  water. 
The  free  acid  or  the  anhydride  has  not  been  prepared.  The  manganous  salt 
is  prepared  by  the  action  of  a  solution  of  sulphurous  acid  upon  manganese 
dioxide  at  a  low  temperature:  Mn02  +  2H2S03  =  MnS,O6  +  2H20  .  Similar 
results  are  obtained  with  nickelic  or  ferric  oxides  (Spring  and  Bourgeois,  BL, 
1886,  46,  151).  The  acid  is  obtained  by  treating  the  manganous  salt  with 
Ba(OH)2  and  the  nitrate  from  this  with  the  calculated  amount  of  H2S04  . 
It  is  a  colorless  solution  and  may  be  evaporated  in  a  vacuum  until  it  has  a 
specific  gravity  of  1.347.  It  decomposes  upon  further  heating:  H2S2O6  =  H2S04 
-f  SO,  .  All  other  thionic  compounds  decompose  upon  heating  witli  separation  of 
sulphur .  By  exposure  to  the  air  dithionic  acid  is  oxidized  to  sulphuric  acid. 
All  dithionates  are  soluble  in  water  and  may  be  purified  by  evaporation  and 
crystallization  (Gelis,  A.  Ch.,  1862,  (3),  65,  230). 

Dithionic  acid  is  also  prepared  by  carefully  adding  a  potassium  iodide  solu- 
tion of  iodine  to  sodium  acid  sulphite  (Hoist  and  Otto,  Arch.  Phann.,  1801,  229, 
171);  Spring  and  Bourgeois  (Arch.  Pharm.,  1891,  229,  707)  contradict  the  above 
statement. 


§261.  Trithionic  acid.     H2S306  =  194. 196 . 

0  0 

II  II 

H'2(S3)100-26 ,  H  — 0  — S  — S  — S  — 0  — H. 
II  II 

0  0 

The  free  acid  and  anhydride  are  not  known.  The  potassium  salt  is  prepared 
by  boiling  potassium  acid-sulphite  with  sulphur  (a) ;  by  treating  potassium 
thiosulphate  with  sulphurous  acid  (&)  (no  action  with  sodium  thiosulphate) 
(Baker,  C.  N.,  1877,  36,  203;  Villiers,  C.  r.,  1889,  108,  402);  by  the  action  of 
iodine  on  a  mixture  of  sodium  sulphite  and  thiosulphate  (c)  (Spring,  B.t  1874, 
7,  1157): 

(a)     12KHS03  +  S2  =  4K2S306  +  2K2S03  +  6H2O 

(6)     4K2S203  +  6S02  =  4K,S306  +  S2 

(c)     Na2S03  +  Na2S203  +  I2  =  Na2S306  +  2NaI 

The  acid  is  prepared  by  adding  perchloric  or  fluosilicic  acid  to  the  potassium 
salt.  The  acid  is  quite  unstable;  at  low  temperature  in  a  vacuum  it  decom- 
poses into  S02  ,  S  and  H2SO4  .  The  salts  are  quite  stable;  they  are  not  oxidized 
by  chloric  or  iodic  acids,  while  the  free  acid  is  rapidly  oxidized  by  these  acids. 
Fixed  alkalis  or  sodium  amalgam  change  the  trithionate  to  sulphite  and  thio- 
sulphate (Spring,  I.e.). 


§263.  TETRATHIONIC  ACID—PENTATHIONIC  ACID.  325 

§262.  Tetrathionic  acid.    H2S,06  =  226.256. 

0  0 

II  II 

H'2(S4)100-26 ,  H  — 0  — S  — S  — S  — S  — 0  — H. 

II  II 

0  0 

The  salts  are  soluble  in  water  and  are  comparatively  stable.  They  are  best 
obtained  in  crystalline  form  by  adding"  alcohol  to  their  solutions  in  water. 
The  acid  has  not  been  isolated  but  it  is  much  more  stable  than  the  tri  or 
pentathionic  acids.  In  dilute  solution  it  can  be  boiled  without  decomposition. 
The  concentrated  solution  decomposes  into  H2S04  ,  S02  and  S  . 

Tetrathionates  are  prepared  by  adding  iodine  to  the  thiosulphates:  2BaS203  + 
I2  =  BaS40Q  +  BaI2  (Maumene,  C.  r.,  1879,  89,  422).  The  lead  salt  is  obtained 
by  the  oxidation  of  lead  thiosulphate  by  lead  peroxide  in  presence  of  sulphuric 
acid:  2PbS2O3  +  PbO2  +  2H2S04  =  PbS406  +  2PbSO4  +  2H20  (Chancel  and 
Diacon,  <7.  pr.,  1863,  90,  55).  To  obtain  the  acid  the  lead  should  be  removed 
by  the  necessary  amount  of  sulphuric  acid,  and  not  by  hydrosulphuric  acid, 
which  causes  the  formation  of  some  pentathionic  acid.  A  number  of  other 
oxidizing-  agents  may  be  used  to  form  the  tetrathionate  from  the  thiosulphate 
(Fordos  and  Gelis,  C.  r.,  1842,  15,  920).  Sodium  amalgam  reconverts  the  tetra- 
thionate into  the  thiosulphate:  Na2S4Os  +  2Na  =  2Na2S203  (Lewes,  /.  (7.,  1880, 
39,  68;  1881,  41,  300).  Tetrathionic  acid  is  also  formed  with  pentathionic  acid 
in  the  reactions  between  solutions  of  H2S  and  SO2  (Wackenroder's  solution, 
A.,  1846,  60,  189).  See  also  Curtius  and  Henkel  (J.  #r.,  1888,  (2),  37,  137).  The 
acid  gives  no  precipitate  of  sulphur  when  treated  with  potassium  hydroxide 
(distinction  from  pentathionic  acid). 

§263.  Pentathionic  acid.     H2S506  =  258.316  . 

0  0 

II  II 

H'2(S5)100-26 ,  H  — 0  — S  — S  — S  — S  — S  — 0  — H. 
II  II 

0  0 

Only  known  in  the  salts  and  in  the  solution  of  the  acid  in  water.  It  is  formed 
by  the  action  of  H2S  upon  S02  in  the  presence  of  water  (a) ;  by  the  action  of 
water  on  sulphur  chloride  (6);  by  the  decomposition  of  lead  thiosulphate  with 
H2S  (Persoz,  Pogg.,  1865,  124,  257) : 

a.     10H2S03  +  10H2S  =  2H2S506  +  5S2  +  18H2O 
6.     10S2C12  +  12H20  =  2H2S506  +  5S2  +  20HC1 

The  filtrate  from  the  decomposition  of  SO,  by  H2S  is  known  as  Wackenroder's 
solution  (Arch.  P/mrw.,  1826,  48,  140).  It  has  been  shown  to  contain  the  tri 
and  tetrathionic  acids  in  addition  to  the  pentathionic  acid  (Debus,  C.  N.,  1888, 
57,  87).  Pentathionic  acid  may  be  concentrated  in  a  vacuum  until  it  has  a 
specific  gravity  of  1.6;  farther  concentration  or  boiling  heat  alone  decomposes 
it  into  H2SO4  ,  SO2  and  S  .  The  solution  of  the  acid  does  not  bleach  indigo. 
When  treated  with  a  fixed  alkali  hydroxide  an  immediate  precipitate  of  sulphur 
is  obtained  (distinction  from  H.,S4O0):  4H-,S,O6  +  20NaOH  =  <>Na2SO3  + 
4NaoS.,O3  +  3S,  +  14H,O  (Takamatsu  and  Smith,  J.  f'.,  1880,  37,  592);  or  if  the 
NaOH  be  added  short  of  neutralization:  10H.,S:.O6  +  20NaOH  =  10Na2S4O.  + 
5S2  +  20H20  .  Neutralization  of  pentathionic  acid  with  barium  carbonate  gives 
barium  tetrathionate  and  sulphur  (Takamatsu  and  Smith,  J.  C.,  1882,  41,  162; 
Lewes,  J,  C.,  1881,  39,  68),  See  also  Spring,  A.,  1879,  199,  97. 


TABLE  OF  THIONIC  ACIDS. 


§264. 


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§265,  4.  SULPHUROUS   ACID.  327 

§265.  Sulphurous  anhydride.    S02  =  64.06. 
Sulphurous  acid.    H2S03  =  82.076. 

0 

II 
SIV0-"2  and  H'2SIV0-"S  ,0  —  S  =  0  and  H  —  0  —  S  —  0  —  H. 

1.  Properties.  —  Sulphurous  anhydride,  S02  ,  sulphur  dioxide,  is  a  colorless  gas 
of  a  strong  suffocating'  odor  of  burning  sulphur.  Specific  gravity  of  the  liquid 
at  0°,  1.4338  (Cailletet  and  Matthias,  C.  r.,  1887,  104,  1563)  ;  of  the  gas  at  0°  and 
760  mm.  pressure,  2.2369  (Leduc,  C.  r.,  1893,  117,  219).  It  is  liquefied  at  atmos- 


white wooly  solid.  Cooled  to  —  76.1°  it  becomes  a  snow-white  solid  (Faraday, 
C.  r.,  1861,  53,  846).  The  dry  gas  is  not  combustible  in  the  air,  does  not  react 
acid  to  litmus,  but  in  presence  of  water  it  has  a  marked  acid  reaction.  The  gas  and 
the  free  acid,  not  the  salts,  are  quite  poisonous,  due  to  the  absorption  of  the 
S02  by  the  blood  and  oxidation  to  H2S04  .  The  gas  is  soluble  in  water,  form- 
ing probably  sulphurous  acid,  H2SO3  .  The  pure  acid  has  not  been  isolated, 
but  forms  salts  mono  and  dibasic  as  if  derived  from  such  an  acid  (Michaelis 
and  Wagner,  B.,  1874,  7,  1073).  It  has  a  strong  odor  from  vaporization  of 
sulphurous  anhydride,  which  is  soon  completely  expelled  upon  boiling.  The 
acid  oxidizes  slowly  in  the  air,  forming  H2SO4  ,  hence  sulphurous  acid  usually 
gives  reactions  for  sulphuric  acid.  Light  seems  to  play  an  important  part  in 
this  oxidation  (Loew,  Am.  8.,  1870,  99,  368).  The  moist  gas  or  a  solution  of  the 
acid  is  a  strong  bleaching  agent,  however  not  acting  alike  in  all  cases.  Wool, 
silk,  feathers,  sponge,  etc.,  are  permanently  bleached:  also  many  vegetable  sub- 
stances, straw,  wood,  etc.;  yellow  colors  and  chlorophyll  are  not  bleached;  red 
roses  are  temporarily  bleached,  immersion  in  dilute  H2S04  restoring  the  color. 

2.  Occurrence.  —  Found  free  in  volcanic  gases  (Ricciardi,  B.,  1887,  20,  464). 

3.  Formation.  —  (a)  By  burning  sulphur  in  air.     (ft)  By  heating  sulphur  with 
various  metallic  oxides,     (c)  By  decomposition  of  thiosulphates  with  HC1.     (rf) 
By  burning  H2S  or  CS2  in  air.     (e)  By  the  action  of  hot  concentrated  sulphuric 
acid  on  metals,  carbon,  sulphur,  etc.     (f)  By  heating  sulphur  with  sulphates. 
(fir)  By  decomposition  of  sulphites  with  acids: 

(a)     S2  +  202  =  2S02 

(&)     Mn02  +  S2  =  MnS  +  SO2 

2Pb304  +  5S2  —  6PbS  +  4S02 

(c)  2Na2S203  -f  4HC1  =  4NaCl  +  2S02  +  S2  +  2H20 

(d)  2H2S  +  302  =  2S02  +  2H20 
CS2  +  3O2  =  2SO,  +  CO, 

(e)  Cu  +  2H2S04  =  CuS04  +  S02  -f  2H2O 
S2  +  4H2S04  =  6S02  +  4H20 

C  +  2H2S04  =  2SO,  +  C02  +  2H2O 
(f  )     FeS04  +  S2  =  FeS  +  2S02 
(g)     Na2S03  +  2H2SO4  =  2NaHSO4  +  SO2  -f  H2O 

4  Preparation.  —  (a)  By  heating  moderately  concentrated  sulphuric  acid  with 
copper  turnings:  Cu  +  2H2SO4  =  CuSO4  -f  SO2  +  2H2O  .  The  gas  is  dried  by 
passing  through  concentrated  sulphuric  acid.  (&)  By  heating  a  mixture  of 
sulphur  and  cupric  oxide  in  a  hard  glass  tube,  (c)  In  a  Kipp's  generator  by 
decomposing  cubes  composed  of  three  parts  calcium  sulphite  and  one  part  of 
calcium  sulphate,  with  dilute  sulphuric  acid  (Neumann,  B.,  1887,  20,  1584). 

Preparation  of  sulphites.—  The  sulphites  of  the  ordinary  metals  are  usually 
made  by  action  of  sulphurous  acid  upon  the  oxides  or  hydroxides  of  the  metals. 
are  normal,  except  mercurous,  which  is  acid,  ancl  chromium,  aluminum 


328  SULPHUROUS  ACID.  §265,  5. 

and  copper,  which  are  basic.     Sulphurous  acid  precipitates  solutions  of  metals 
of  the  first  and  second  groups,  except  copper  and  cadmium. 

The  sulphites  of  the  alkalis  precipitate  solutions  of  the  other  metals  except 
chromium  salts;  and  some  normal  sulphites  may  be  made  in  this  manner. 
The  sulphites  of  silver,  mercury,  copper  and  ferricum  (known  only  in  solution) 
are  unstable,  the  sulphurous  acid  becoming  sulphuric  at  the  expense  of  the 
base,  which  is  reduced  to  a  form  having  a  less  number  of  bonds.  With  the 
unstable  stannous  sulphite  the  action  is  the  reverse.  (See  6 A.)  All  sulphites  by  ex- 
posure to  the  air  slowly  absorb  oxygen,  and  are  partially  converted  into  sulphates. 

5.  Solubilities. — One  volume  of  water  at  0°  dissolves  68.861  volumes  of  sul- 
phurous anhydride;  at  20°,  36.206  volumes  (Carius,  A.,  1855,  94,  148);   or  at  20°  , 
0.104  part  by  weight  (Sims  J.  C.,  1862,  14,  1).     One  volume  of  alcohol  dissolves 
at   15°,    116  vol.   SO2 .     Charcoal   absorbs  165  volumes,  camphor  308  volumes, 
glacial  acetic  acid  318  volumes  of  the  gas.     Liquid  sulphurous   anhydride   dis- 
solves P  ,  S  ,  I ,  Br  and  many  gases. 

The  sulphites  of  the  metals  of  the  alkalis  are  freely  soluble  in  water;  the 
normal  sulphites  of  all  other  metals  are  insoluble,  or  but  very  slightly  soluble 
in  water.  The  sulphites  of  the  metals  of  the  alkaline  earths,  and  some  others, 
are  soluble  in  solution  of  sulphurous  acid,  the  solution  being  precipitated  on 
boiling.  The  alkali  bases  form  acid  sulphites  (bisulphites),  which  can  be 
obtained  in  the  solid  state,  but  evolve  sulphurous  anhydride.  The  sulphites 
are  insoluble  in  alcohol.  They  are  decomposed  by  all  acids  except  carbonic 
and  boric,  and  in  some  instances,  hydrosulphuric. 

6.  Reactions.     A— With  metals  and  their  compounds. — Sulphurous  acid 
reacts  with  Zn ,  Fe ,  Sn ,  and  Cu  to  form  hyposulphurous  acid,  H2SO, 
(Schiitzenberger,  C.  r.,  1869,  69,  196).     With  Zn  in  the  presence  of  HC1 
it  is  reduced  to  hydrosulphuric  acid:  3Zn  +  6HC1  +  H2S03  =  3ZnCl2  + 
H2S  -f-  3H20  .     Free  sulphurous  acid  precipitates  solutions  of  first  and 
second  group  metals  except  those  of  copper  and  cadmium;  solutions  of 
other  metallic  salts  are  not  precipitated  owing  to  the  solubility  of  the 
sulphites  in  acids. 

Alkali  sulphites  precipitate  solutions  of  all  other  metallic  salts.  The 
precipitates,  mostly  white,  are  soluble  in  acetic  acid.  The  precipitates 
of  Pb ,  Hg ,  Ba ,  Sr ,  and  Ca  are  usually  accompanied  by  sulphates,  due  to 
the  fact  that  soluble  sulphites  nearly  always  contain  sulphates  (4). 

Solution  of  lead  acetate  precipitates,  from  solutions  of  sulphites,  lead 
sulphite,  PbS03 ,  white,  easily  soluble  in  dilute  nitric  acid ;  and  not  blacken- 
ing when  boiled  (distinction  from  thiosulphate).  Solution  of  silver  nitrate 
gives  a  white  precipitate  of  silver  sulphite,  Ag2S03 ,  easily  soluble  in  very 
dilute  nitric  acid  or  in  excess  of  alkaline  sulphite,  and  turning  dark- 
brown  when  boiled,  by  formation  of  metallic  silver  and  sulphuric  acid. 
Solution  of  mercurous  nitrate  with  sodium  sulphite  gives  a  gray  precipi- 
tate of  metallic  mercury.  Solution  of  mercuric  chloride  produces  no 
change  in  the  cold;  but  on  boiling,  the  white  mercurous  chloride  is  precipi- 
tated, with  formation  of  sulphuric  acid.  Still  further  digestion,  with 
sufficient  sulphite,  reduces  the  white  mercurous  chloride  to  gray  metallic 
mercury  (§58,  6e). 

Solution  of  ferric  chloride  gives  a  red  solution  of  ferric  sulphite, 
Fej(SOg)a  ;  or,  in  more  concentrated  solutions,  a  yellowish  precipitate  of 


§265,  6//,  2.  SULPHUROUS  ACID.  329 

basic  ferric  sulphite,  also  formed  by  addition  of  alcohol  to  the  red  solu- 
tion. The  red  solution  is  decolored  on  boiling;  the  acid  radical  reducing 
the  basic  radical,  and  forming  ferrous  sulphate. 

Solution  of  barium  chloride  gives  a  white  precipitate  of  barium  sul- 
BaSO;!  ,  easily  soluble  in  dilute  hydrochloric  acid — distinction  from 
,  which  is  imdis^olved,  and  should  be  filtered  out.  Now,  on  adding 
to  the  filtrate  nitrohydrochloric  acid,  a  precipitate  of  barium  sulphate 
is  obtained — evidence  that  sulphite  has  been  dissolved  by  the  hydrochloric 
acid: 

BaS03  +  2HC1  =  BaCL  +  H2SO;! 

BaCl2  +  H2S03  +  C12  +  H2O  =  BaS04  +  4HC1 

One  part  of  barium  sulphite  is  dissolved  by  46,000  parts  of  water 
at  18°. 

Calcium  chloride  reacts  similar  to  barium  chloride,  the  precipitate  of 
calcium  sulphite  being  less  soluble  in  water  than  the  corresponding  sul- 
phate. One  part  of  calcium  sulphite  is  dissolved  by  800  parts  of  water 
at  18°  while  one  part  of  the  strontium  salt  is  dissolved  by  30,000  parts 
water  at  18°. 

Sulphurous  acid  and  sulphites  are  active  reducing  agents  by  virtue  of 
their  capacity  for  oxidation  to  sulphuric  acid  and  sulphates. 

The  reactions  with  silver,  mercury  and  ferricum  given  above  illustrate 
the  reducing  action,  and  the  following  should  also  be  noted: 

Pb02  becomes  lead  sulphate. 

Asv  forms  arsenous  and  sulphuric  acids. 

Sbv  forms  Sb'". 

Cu"  becomes  cuprous  sulphate. 

CrVI  forms  chromic  sulphate. 

Co'"  forms  cobaltous  sulphate. 

Ni'"  forms  nickel  sulphate. 

Mn"+n  forms  manganous  sulphate. 

With  Mn02  in  the  cold,  manganous  dithionate,  MnS206 ,  is  formed 
(Gmelin's  Hand-look.,  2,  174). 

With  stannous  chloride  sulphurous  acid  acts  as  an  oxidizing  agent,  form- 
ing stannic  sulphide  and  stannic  chloride  or  stannic  chloride  and  hydro- 
sulphuric  acid,  according  to  the  amount  of  hydrochloric  acid  present 
(§71,  <fe). 

B. — With  non-metals  and  their  compounds. — Upon  other  acids  sul- 
phurous acid  acts  as  a  reducing  agent,  except  with  hypophosphorous,  phos- 
phorous, and  hydrosulphuric  acids. 

1.  H,Fe(CN)6  forms  H4Fe(CN)6  and  H2S04 . 

2.  HN02  and  HN03  form  NO  and  H2SO< . 


330  SULHPUROUS   ACID.  §«65,  6£,  3. 

3.  PH3  +  2H2S03  =  H,P04  +  S2  +  2H20  (Carvazzi,  Gazzetta,  1886,  16, 
169).     H3P02  becomes  H3P04  and  the  S02  is  reduced  to  S ,  and  with  excess 
of  H3P02"to  H2S .     H3PO,.  forms  H,P04  and  H2S  ($253,  6). 

4.  H2S  forms  S  from  both  compounds :   4H2S  +  2S02  =  3S2  +  4H&0  . 
See  also  §263  . 

5.  Cl ,  HC10 ,  and  HC103  form  hydrochloric  and  sulphuric  acids. 

6.  Br    forms    hydrobromic    and    sulphuric    acids.     HBr03    forms    first 
bromine  then  hydrobromic  acid,  sulphuric  acid  in  both  cases. 

7.  I  forms  hydriodic  and  sulphuric  acids.     Tn  presence  of  hydrochloric 
acid   and   a   barium   salt   it   serves   as   a   means   of   detecting   a   sulphite 
mixed  with  a  sulphate  and  a  thiosulphate  (Smith,  C.  N.,  1895,  72,  39). 
HI03  forms  first  iodine  then  hydriodic  acid,  sulphuric  acid  in  both  cases. 

7.  Ignition. — Acid  sulphites  heated  in  sealed  tube  to  150°  are  decomposed 
into  sulphates  and  sulphur  (Barbaglia  and  Gucci,  J?.,  1880,  13,  2325;  Berthelot, 
A.  Ch.,  1864,  (4),  1,  392).  Dry  SO2  at  high  heat  with  many  metals  is  decom- 
posed, forming  a  sulphide  and  sulphate  or  sulphite  (Uhl,  B.,  ls<)0,  23,  :2i:>l). 
Sulphites  are  decomposed  by  heat  into  oxides  and  sulphurous  anhydride: 
CaSO3  —  CaO  +  SO,;  or  into  sulphates  and  sulphides:  4NaoS03  =  3Na.,S04  + 
Na2S. 

8.  Detection. — Free  sulphurous  acid  is  detected  by  its  odor  and  by  its 
decolorizing  action  upon  a  solution  of  KMn04  or  I  (Hilger,  J.  (7.,  1876, 
29,  443).     The  reaction  with  iodic  acid  is  also  employed  as  a  test  for 
sulphurous  acid  (as  well  as  for  iodic).     A  mixture  of  iodic  acid  and  starch 
is  turned  violet  to  blue  by  traces  of  sulphurous  acid  or  sulphites  in  vapor 
or  in  solution,  the  color  being  destroyed  by  excess  of  the  sulphurous  acid 
or  the  sulphite.     Sulphites  are  distinguished  from  sulphates  by  failure  to 
precipitate  with  BaCl2  in  presence  of  HC1 .     After  removal  of  the  BaS04 
by  filtration  the  sulphite  is  oxidized  to  sulphate  by  chlorine  water  and 
precipitated  by  the  excess  of  BaCl2  present.    For  separation  from  sulphides 
and  thiosulphates  see  §258,  8. 

Normal  potassium  sulphite,  K2S03 ,  is  alkaline  to  litmus  but  when 
treated  with  BaCl2  gives  a  neutral  solution.  The  acid  sulphite,  KHS03 , 
is  neutral  to  litmus  but  with  BaCl2  gives  an  acid  solution:  2KHSOn  + 
BaCL  =  BaS03  +  2KC1  +  S02  '+  H26  (Villiers,  C.  r.,  1887,  104,  1177). 

9.  Estimation. —  (rt)  After  converting  into  H2S04  by  HN03  or  Cl  it  is  precipi- 
tated by  BaCL  and  weighed  as  BaSO4  .  (6)  The  oxidation  is  effected  by  fusing 
with  Na2C03  and  KNO3  (equal  parts),  (c)  A  standard  solution  of  iodine  is 
added,  and  the  excess  of  iodine  determined  by  a  standard  solution  of  N"a3S203  . 


§266,  4.  SULPHURIC  ACID. 

§266.  Sulphuric  acid.     H6S04  =  98.076. 

0 

H',SVI0~"4 ,  H  -  0       S  —  0  —  H  . 

II 
0 

1.  Properties. — Absolute    sulphuric    acid,    H2S04  ,   is    a    colorless    oilv    liquid 
(oil  of  vitriol):  spec-Hie  gravity,   l.S.'iTl  at  15°    (Meiidelejeff,  #.,  1884,   17,   2541). 
According1  to   Marignac   (.1.   1'Ji..  isr>:{.   (:;),  39,  184),  it  begins  to  boil  at  about 
290°,   ascending'   to    :-i:!S°    with    partial    decomposition.     At   temperatures   much 
below  the  boiling-  point  (160°)  it  vaporizes  from  open  vessels,  giving  off  heavy, 
white,   suffocating   vapors,   exciting-  coughing-  without   giving  premonition   by 
odor.     At  ordinary  temperature  it  is  non-volatile  and  inodorous.     At  low  tem- 
peratures it   solidifies   to    a   crystalline   mass.     The   freezing   point  is   greatly 
influenced  by  the  amount  of  water  present.     When  the  acid  contains  one  mole- 
cule of  water,   H2SO4.H2O  ,   the  melting  point  is  highest,    -f-7.5°    (Pierre   and 
Puchot,  A.  Ch.,  1874,  (5),  164). 

H2S04  is  a  very  strong  acid  and,  because  of  its  high  boiling  point, 
displaces  all  the  volatile  inorganic  acids;  on  the  other  hand  it  is  displaced, 
when  heated  above  its  boiling  point,  by  phosphoric,  boric,  and  silicic  acids. 
It  is  a  dibasic  acid,  forming  two  series  of  salts,  M'HS04  and  M'2S04 .  It  is 
miscible  with  water  in  all  proportions  with  production  of  heat;  it  abstracts 
water  from  the  air  (use  in  desiccators),  and  quickly  abstracts  the  elements 
of  water  from  many  organic  compounds,  and  leaves  their  carbon,  a  char- 
acteristic charring  effect.  It  dissolves  in  alcohol,  without  decomposing  it 
—but  if  in  sufficient  proportion  producing  ethylsulphuric  acid,  HC2H5S04 . 

Sulphuric  anhydride,  S03  ,  is  a  colorless,  fibrous  or  waxy  solid,  melting-  at 
14.8°  (Rebs,  A.,  1888,  246,  379),  boiling-  at  46°  (Schulz-Sellak,  B.,  1870,  3,  215), 
and  vaporizing  with  heavy  white  fumes  in  the  air  at  ordinary  temperatures. 
It  is  very  deliquescent,  and  on  contact  with  water  combines  rapidly,  forming 
sulphuric  acid  with  generation  of  much  heat. 

2.  Occurrence. — Found  free  in  the  spring  water  of  volcanic  districts.     Found 
combined  in  g-ypsum,  CaS04  +  2HoO;  in  heavy  spar,  BaS04;  in  celestite,    SrSO4; 
in  Epsom  salts,  MgS04  +  7H20;  in  Glauber  salt,  Na2S04  +  10H2O  ,  etc. 

3.  Formation. —  (a)  By  electrolyzing-  H,O  ,  using  Pt  electrodes  with  pieces  of 
S  attached  (Becquerel,  C.  r.,  1863,  56,  237)".   (6)  By  oxidizing  S  or  SO2  in  presence 
of  water  by  Cl ,  Br  ,  HNO3  ,  etc.     (c)  By  heating  S  and  H20  to  200°.     (d)   By 
adding  H26  to  S03.     (e)   By  passing-  a  mixture   of  S02   and  0   over  platinum 
spong-e  and  then  adding-  water. 

4.  Preparation. — Industrially,  sulphuric  acid  is  made  by  utilizing  the 
S02  evolved  as  a  by-product  in  roasting  various  sulphides — e.  g.,  iron  and 
copper  pyrites,  blende,  etc.  (a)  and  (&);  or  by  burning  sulphur  in  the  air 
to  form  the  S02 .  The  S02  is  oxidized  and  converted  into  sulphuric  acid 
by  two  distinct  processes  known  as  the  contact  and  the  chamber  process. 
In  the  contact  process  the  S02 ,  after  careful  purification,  arsenic  especially 
being  removed,  is  passed  together  with  oxygen  through  a  contact  mass 


33$  SULPHURIC  ACID.  §266,  4. 

containing  finely  divided  platinum  or  other  catalytic  reagent  maintained 
at  the  proper  temperature.  The  S02  unites  with  the  oxygen  to  form 
S03 ,  which  -is  absorbed  in  dilute  sulphuric  acid.  This  process  is  espe- 
cially advantageous  for  making  concentrated  or  fuming  sulphuric  acid. 
In  the  chamber  process  the  S02  is  passed  into  a  large  leaden  chamber  and 
brought  into  contact  with  HN03 ,  steam,  and  air.  The  HN03  first  oxidizes 
a  portion  of  the  S02  (c)  ;  the  steam  then  reacts  upon  the  NO. ,  forming 
HN03  and  NO  (d).  This  NO  is  at  once  oxidized  again  by  the  air  to  N02 , 
so  that  theoretically  no  nitric  acid  is  lost,  but  all  is  used  over  again. 
Practically,  traces  of  it  are  constantly  escaping  with  the  nitrogen  intro- 
duced as  air,  so  that  a  fresh  supply  of  nitric  acid  is  needed  to  make  up  for 
this  loss.  The  dilute  acid  known  as  chamber  acid  is  concentrated  first 
in  lead  pans  and  then  in  platinum  or  silica  pans.  Commercial  sulphuric 
acid  known  as  oil  of  vitriol  has  a  sp.  gr.  of  1.83  and  contains  93% 
K,S04  ;  when  heated  to  338°  a  98%  acid  distills  over.  The  absolute 
H2S04  cannot  be  made  by  evaporation  or  distillation;  it  still  contains 
about  two  per  cent  of  water.  It  may  be  made  by  adding  to  water,  or  to 
the  H2S04  containing  the  two  per  cent  of  water,  a  little  more  S03  or 
H2S207  than  would  be  needed  to  make  H2S04  ;  then  passing  perfectly 
dry  air  through  it  until  the  excess  of  S03  is  removed,  leaving  absolute 
H2S04 .  Fuming  pyrosulphuric,  or  Nordhausen  sulphuric  acid,  H2S207 , 
is  made  by  solution  of  sulphuric  anhydride  in  sulphuric  acid  (e)  ;  by 
drying  FeS04  +  7H,0  until  it  becomes  FeS04  +  H20  ,  and  then  distilling 
(/).  Sulphuric  anhydride  is  made  by  the  action  of  heat  on  sodium 
pyrosulphate,  Na2S207  (g),  prepared  by  heating  NaHS04  to  dull  redness ;  by 
distilling  pyrosulphuric  acid,  the  anhydride  is  collected  in  an  ice-cooled 
receiver;  by  heating  H2S04  with  P205  (7i): 

(a)     2ZnS  +  302  =  2ZnO  +  2S02 

(&)     4FeS2  +  11O3  =  2Fe203  -f  SSO3 

(c)  S02  +  2HNOS  =  H2S04  +  2N02 

(d)  3N02  +  H20  =  2HN03  +  NO 

(e)  H2S04  +  S03  =  H2S207 

(0     4FeS04  +  H20  =  2Fe203  +  H2S207  +  2SO, 

(g)     Na2S207  =  Na2S04  +  S03 

(ft)     H2S04  +  P205  =  2HP03  +  S03 

Sulphates  are  made:  (a)  by  dissolving  the  metals  in  sulphuric  acid; 
(#)  by  dissolving  the  oxides  or  hydroxides;  (c)  by  displacement.  All 
salts  containing  volatile  acids  are  displaced  by  sulphuric  acid  and  a 
sulphate  formed  (except  the  chlorides  of  mercury).  The  excess  of  acid 
may  generally  be  expelled  by  evaporation,  or  the  cry  stills  washed  with 
cold  water  or  alcohol.  The  insoluble  sulphates  are  best  made  by  precipita- 
tion. 


§266,  6A.  SULPHURIC  ACID.  '    333 

5.  Solubilities. — Sulphuric  acid  is  miscible  with  water  in  all  proportions; 
the  concentrated  acid  with  generation  of  much  heat.  Sulphuric  acid 
transposes  the  salts  of  nearly  all  other  acids,  forming  sulphates,  and  either 
acids  (as  hydrochloric  acid,  $269,  4)  or  the  products  of  their  decomposi- 
tion (  as  with  chloric  acid,  §273,  6).  Chl<>i  <  of  silver,  tin,  and  antimony 
are  with  difficulty  transposed  by  sulphuric  acid,  and  chlorides  of  mercury 
not  at  all.  Also,  at  temperatures  above  about  300°  phosphoric  and  silicic 
acids  (and  other  acids  not  volatile  at  this  temperature)  transpose  sulphates, 
with  vaporization  of  sulphuric  acid. 

The  sulphates  of  Pb ,  Hg',  Ba,  Sr,  and  Ca  are  insoluble,  those  of  Hg' 
and  Ca  sparingly  soluble.  Sulphuric  acid  and  soluble  sulphates  precipi- 
tate solutions  of  the  salts  of  Pb ,  Hg',  Ba ,  Sr ,  and  Ca  ;  Hg'  and  Ca  salts 
incompletely.  The  metallic  sulphates  are  insoluble  in  alcohol  which  pre- 
cipitates them  from  their  moderately  concentrated  aqueous  solutions. 
Alcohol  added  to  solutions  of  the  acid  sulphates  precipitates  the  normal 
sulphates,  sulphuric  acid  remaining  in  solution:  2KHS04  =  K2S04  + 
H2S04 .  PbS04  is  soluble  in  a  saturated  solution  of  NaCl  in  the  cold, 
depositing  after  some  time  crystals  of  PbCl2 ,  complete  transposition  being 
effected.  A  solution  of  PbCl2  in  NaCl  is  not  precipitated  on  addition  of 
H2S04  (Field,  J.  C.,  1872,  25,  575). 

6.  Reactions.  A. — With  metals  and  their  compounds. — Sulphuric  acid, 
dilute,  has  no  action  on  Pb  ,  Hg ,  Ag ,  Cu  *,  and  Bi .  Au  ,  Pt ,  Ir ,  and  Kh 
are  not  attacked  by  the  acid,  dilute  or  concentrated;  other  metals  are 
attacked  by  the  hot  concentrated  acid  with  evolution  of  S02 .  The  fol- 
lowing metals :  Sn ,  Th ,  Cd ,  Al ,  Fe ,  Co ,  Ni ,  Mn ,  Zn  ,  Mg ,  K  ,  and  Ka 
are  attacked  by  the  acid  of  all  degrees  of  concentration ;  the  dilute  rapidly 
and  the  cold  concentrated  slowly,  with  evolution  of  hydrogen;  the  hot 
concentrated  with  evolution  of  S02 .  The  degree  of  concentration  and 
the  temperature  may  be  regulated  so  that  the  two  gases  may  be  evolved 
in  almost  any  desired  proportions.  A  secondary  reaction  frequently  takes 
place,  the  metal  decomposing  the  S02  forming  H2S  or  a  sulphide;  and  the 
H2S  decomposing  the  S02  with  separation  of  sulphur  (Ditte,  A.  Ch.,  1890, 
(6),  19,  68;  Muir  and  Adie,  J.  C.,  1888,  53,  47). 

Sulphuric  acid  or  soluble  sulphates  react  with  soluble  barium  salts  to 
give  barium  sulphate,  white,  insoluble  in  hydrochloric  or  nitric  acids.  This 
insolubility  is  a  distinction  from  all  other  acids  except  selenic  and  fluo- 
silicic.  The  precipitate  formed  in  the  cold  is  very  fine  and  difficult  to 
separate  by  nitration;  if  formed  in  hot  acid  solution  and  then  boiled  it  is 
retained  by  a  good  filter.  In  dilute  solution  for  complete  precipitation 
the  mixture  should  stand  for  some  time.  Solutions  of  lead  salts  give  a 

*  Andrews,  J.  Am.  Soc,,  1896, 18,  251. 


334  SULPHURIC  ACID.  §266,  6B. 

white  precipitate  of  lead  sulphate  not  transposed  by  acids  except  H2S  (5), 
soluble  in  the  fixed  alkalies.  The  presence  of  alcohol  makes  the  precipi- 
tation quantitative  (§57,  9).  Solution  of  calcium  salts  not  too  dilute  form 
a  white  precipitate  of  calcium  sulphate  (§188,  oc). 

Dilute  sulphuric  acid  does  not  oxidize  any  of  the  lower  metallic  oxides. 
Concentrated  sulphuric  acid  is  an  oxidizing  agent.  \Yhen  hot  it  liberates 
one  atom  of  oxygen  and  is  reduced  to  sulphurous  acid,  which  is  decom- 
posed with  the  evolution  of  sulphur  dioxide  and  water. 

The  concentrated  acid  with  the  aid  of  heat  effects  the  following 
changes: 

Hg20  forms  mercuric  sulphate,  and  sulphurous  anhydride  is  evolved. 

SnCl2  forms,  first,  sulphurous  anhydride,  then  hydrosulphuric  acid, 
stannic  chloride  at  the  same  time  being  produced. 

Fe"  is  changed  to  Fe2(S04)a  by  hot  concentrated  sulphuric  acid. 

Mn"+n  forms  MnS04  and  0 .  That  is,  all  compounds  of  manganese 
having  a  degree  of  oxidation  above  the  dyad  are  reduced  to  the  dyad  with 
evolution  of  oxygen. 

Potassium  permanganate  dissolves  in  cold  concentrated  sulphuric  acid 
with  formation  of  a  green  solution  of  a  sulphate  of  the  heptad  manganese, 
(Mn03)2S04  (§134,  5c). 

Similarly  the  hot  concentrated  acid  also  reduces  Pb1^  to  Pb",  Co'"  to 
Co",  Ni'"  to  Ni",  FeVI  to  Fe'",  and  CrVI  to  Cr'",  oxygen  being  liberated 
(oxidized)  and  the  metal  reduced  while  the  bonds  of  the  S04  radical  are 
not  changed;  a  sulphate  of  the  metal  being  produced. 

B. — With  non-metals  and  their  compounds. — When  dilute  sulphuric  acid 
transposes  the  salts  of  other  acids,  no  other  change  occurs  if  the  acid  set 
free  be  stable  under  the  conditions  of  its  liberation.  In  ordinary  reactions 
sulphuric  acid  never  acts  as  a  reducing  agent. 

1.  Many  organic  acids  and  other  organic  compounds  are  decomposed  by 
the  hot  concentrated  acid,  the  elements  of  water  being  abstracted  and 
carbon  set  free.     Continued  heating  of  the  carbon  with  the  hot  concen- 
trated acid  oxidizes  it  to  C02  with  liberation  of  S02 . 

H2C204  becomes  C02 ,  CO ,  and  H20  .  The  bonds  of  the  H2S04  remain 
unchanged. 

K4Fe(CN)6  with  dilute  H2S04  forms  HCN  :    2K4Fe(CN)G  +  3H0S04  = 
6HCN  +  K2FeFe(CN),  +  3K2S04 . 

Cyanates  are  decomposed  into  C02  and  NH3 :  2KCNO  +  2H2S04  +  2H20 
=  K2S04  +  (NH4)2S04  +  2C02 . 

Thiocyanates  are  also  decomposed  by  concentrated  sulphuric  acid. 

2.  Nitrites   are    decomposed   with   formation   of   nitric   acid  and   NO  : 
6KN02  +  3H2S04  =  3K2S04  +  2HNO:]  +  4NO  +  2H20  . 


$266,  8.  SULPHURIC  ACID.  335 

3.  H3P02  or  hypophosphites  are  oxidized  to  phosphoric  acid  with  re- 
duction of  the  sulphuric  acid  to  sulphurous  acid  and  then  to  sulphur. 

4.  Sulphur  is  slowly  changed  by  hot  concentrated  sulphuric  acid  to 
sulphurous  acid  with  reduction  of  the  sulphuric  acid  to  the  same  com- 
pound.     Hydrosulphuric   acid  with  hot  concentrated  sulphuric   acid  is 
oxidized  to  sulphur  with  reduction  of  the  sulphuric  acid  to  sulphurous 
acid.     Further  oxidation  may  take  place  as  indicated  ahove. 

•5.  Chlorates  are  transposed  and  then  decomposed  when  treated  with 
concentrated  sulphuric  acid :  3KC103  +  2H2S04  ==  2KHS04  +  KC104  + 
2C102  +  HL>0  . 

6.  HBr  forms  Br  and  S02 .     No  action  except  in  concentrated  solution. 

7.  HI  forms  I  and  S02 . 

7.  Ignition. — All   sulphates  fused  with   a  fixed   alkali   carbonate   are 
transposed  to  carbonates  (oxide  or  metal  if  the  carbonate  is  decomposed 
by  the   heat  used,   §228,  7)   with   formation  of   a  fixed   alkali   sulphate 
(method  of  analysis  of  insoluble  sulphates).     If  the  sulphate,  or  any  other 
compound  containing  sulphur,   is  fused  in  the  presence   of  carbon,  as 
fusion  with  a  fixed  alkali  carbonate  on  a  piece  of  charcoal,  the  resulting 
mass  contains  an  alkali  sulphide,  which,  when  moistened,  blackens  metallic 
silver. 

The  sulphates  of  Cu  ,  Sb  ,  Fe  ,  Hg ,  Ni  and  Sn  are  completely  decomposed  at 
a  red  heat:  2FeSO4  =  Fe2O3  +  S03  +  S02;  2CuS04  =  2CuO  +  2S02  -f  O2  .  A 
white  heat  decomposes  the  sulphates  of  Al ,  Cd  ,  Ag  ,  Pb  ,  Mn  and  Zn  .  An 
ordinary  white  heat  has  no  action  on  the  sulphates  of  the  alkalis  and  alkaline 
earths;  but  at  the  most  intense  heat  procurable  the  sulphates  of  Ba ,  Ca  and 
Sr  are  changed  to  oxides;  and  at  the  same  temperature  K2S04  and  Na2SO4  are 
completely  volatilized,  preceded  by  partial  decomposition. 

Lead  sulphate  heated  in  a  current  of  hydrogen  is  reduced  according  to  the 
following  equation:  2PbS04  +  6H2  =  Pb  +  PbS  +  SO2  +  6H,0  .  After  a 
distinct  interval  the  remainder  of  the  sulphur  is  removed  as  H2S:  PbS  -f-  H2  = 
Pb  +  H:S  (Rodwell,  J.  C.,  1863,  16,  42).  Potassium  sulphate  heated  in  a 
current  of  hydrogen  is  reduced  to  potassium  acid-sulphide:  K2S04  -f-  4H-  = 
KOH  +  KHS  +  3H2O  (Berthelot,  A.  Cft,,  1890,  (6),  21,  400).  Potassium  acid- 
sulphate,  KHSO4  ,  heated  to  200°  evolves  H2SO4  .  The  sodium  acid-sulphate 
decomposes  more  readily. 

8.  Detection.  — Free  sulphuric  acid  or  the  soluble  sulphates  are  detected 
by  precipitation  in  hot  hydrochloric  acid  solution  with  barium  chloride, 
forming  the  white,  granular,  insoluble  barium  sulphate. 

The  sulphates  insoluble  in  water  are  decomposed  for  analysis — (1st)  by 
long  boiling  with  solution  of  alkali  carbonate;  and  more  readily  (2d)  by 
fusion  with  an  alkali  carbonate.  In  both  cases  there  are  produced — alkali 
mil jrtid lex  soluble  in  water,  and  carbonates  soluble  by  hydrochloric  or  nitric 
acid,  after  removing  the  sulphate  (a).  If  the  fusion  be  done  on  charcoal, 
more  or  less  dcoxidation  will  occur,  reducing  a  part  or  the  whole  of  the 


336  PERSULPHURIC  ACID.  §266,  9. 

sulphate  to  sulphide  (7),  and  the  carbonate  to  metal  (as  with  lead,  §57,  7), 
or  leaving  the  metal  as  a  carbonate  or  oxide  (7,  §§222  and  228). 

a.  BaS04  +  Na2C03  =  Na2S04  (soluble  in  water)  +  BaC03   (soluble  in  acid). 

A  mixture  of  H,S04  and  a  sulphate  may  be  separated  by  strong  alcohol, 
which  precipitates  the  latter.  A  test  for  free  sulphuric  acid,  in  distinction  from 
sulphates,  may  be  made  by  the  use  of  cane  sug-ar,  as  follows:  A  little  of  the 
liquid  to  be  tested  is  concentrated  on  the  water-bath;  then  from  two  to  four 
drops  of  it  are  taken  on  a  piece  of  porcelain,  with  a  small  fragment  of  white 
sugar,  and  evaporated  to  drj'ness  by  the  water-bath.  A  greenish-black  residue 
indicates  sulphuric  acid.  (With  the  same  treatment,  hydrochloric  acid  gives  a 
brownish-black,  and  nitric  acid  a  yellow-brown  residue.)  A  strip  of  white 
glazed  paper,  wet  with  the  liquid  tested,  by  immersing  it  several  times  at  short 
intervals,  then  dried  in  the  oven  at  100°,  will  be  colored  black,  brown  or  reddish, 
if  the  liquid  contains  as  much  as  0.2  per  cent  of  sulphuric  acid. 

9.  Estimation  __  (a)  By  precipitation  as  barium  sulphate  and  weighing  as 
such.  The  solution  should  be  hot  and  acidified  with  hydrochloric  acid,  and 
the  mixture  should  be  boiled  a  few  minutes  after  the  addition  of  the  barium 
chloride.  (6)  By  precipitation  as  barium  sulphate  with  an  excess  of  an  hydro- 
chloric acid  solution  of  barium  chromate  (three  per  cent  hydrochloric  acid). 
Add  NH4OH  ,  fill  to  a  definite  volume,  and  filter  through  a  dry  filter-paper. 
Transfer  an  aliquot  portion  to  an  azotometer  with  H202  ,  and  after  acidifying, 
determine  the  oxygen  evolved  (Baumann,  Z.  anc/ew.,  1891,  140)  (§244,  6A,  12) 
(e)  When  present  in  small  amounts  in  drinking  water  by  a  photometric  method 
(Hinds,  C.  N.,  1896,  73,  285  and  299). 

§267.  Persulphuric  acid.   HS04  =  97.068. 

1.  The  anhydride.  —  The  anhydride,  S2O7  ,  was  discovered  by  Berthelot  (C.  r., 
1878,  88,    20  and  71).     It  is  obtained  by  the  action  of  the  silent  electric  discharge 
upon  a  mixture  of  equal  volumes  of  dry  SO2  and  O  .     At  0°,  it  consists  of  flexible 
cyrstalline  needles,  remaining  stable  for  several  days.     When  heated  it  decomposes 
into  SO3  and  O  .     With  SO2  it  combines  to  form  SO  3:  S2O7+SO2  =  3SO3  .    Although 
in  its  reactions  it  acts  as  a  strong  oxidizing  agent,  it  is  weaker  than  chlorine  or  ozone; 
oxalic  acid  and  chromium  salts  are  not  oxidized  (Traube,  B.,  1889,  22,  1518,  1528; 
1892,  25,  95). 

2.  The  Acid.  —  The  acid  was  first  prepared  by  Marshall,  who  electrolyzed  cold 
fairly  dilute  sulphuric  acid  (J  .  C.,  £9,  771).     Hydrogen  is  liberated  at  the  cathode 
while  the  HSO4  anions  discharged  at  the  anode  unite  to  form  persulphuric  acid, 
the  following  reaction  taking  place. 


The  acid  may  also  be  formed  by  the  action  of  H2O2  on  concentrated  H2SO4  .  Water 
solutions  of  the  acid  decompose  very  rapidly.  Solutions  of  the  acid,  in  concen- 
trated sulphuric  acid,  are  more  stable. 

3.  Salts.  —  The  potassium  salt,  K2S2OS  ,  is  prepared  by  the  electrolysis  of  a  sat- 
urated solution  of  KHSO4  with  a  current  of  3  to  3.5  amperes.  It  is  a  white  crys- 
talline powder,  which  may  be  recrystallized  from  hot  water  with  almost  no  decom- 
position. Continued  heating  of  the  solution  effects  decomposition.  One  hundred 
parts  of  water  dissolve  0.564  part  of  the  salt  at  0°  and  4.08  parts  at  40°. 

The  ammonium  salt  is  prepared  by  the  electrolysis  of  a  saturated  solution  of 
ammonium  sulphate.  One  hundred  parts  of  water  dissolve  58.2  parts  of  the  salt 
at  0°.  It  can  be  recrystallized  from  water  if  the  solution  not  heated  above  60°. 
It  forms  monoclinic  crystals.  The  dry  salt  is  stable  at  100°.  It  is  used  in  the 
cyanide  process  for  the  recovery  of  gold  (Elbs,  Z.  annew.,  1897,  195).  The  potas- 
sium is  the  least  soluble  of  the  persulphate  salts.  A  solution  of  KoCOs  gives  an 
abundant  crystalline  precipitate  of  K2S2Og  from  a  solution  of  the  ammonium  salts. 
The  barium  salt,  BaS2Os.4H2p  ,  is  fairly  soluble  and  may  be  prepared  by  rubbing 
the  ammonium  salt  with  barium  hydroxide. 


,  1.  CHLORINE.  33? 

3.  Reactions. — All  persulphates  when  dissolved  in  water  are  decomposed  slowly 
in  the  cold  and  more  rapidly  on  heating,  oxygen,  free  sulphuric  acid  and  a  sulphate 
being  formed. 

2K2S2OS  +  2H2O  =  2K2SO4  +  2H2SO4  +  O2 . 

A  large  proportion  of  the  oxygen  escapes  as  ozone,  which  may  be  identified  by  its 
odor  arid  action  on  starch  iodide  paper.  Ammonium  persulphate  in  water  solution 
decomposes  slowly  at  the  ordinary  temperature  without  the  evolution  of  oxygen. 

8(NH4)2S2OS  +  6H2O  =  14NH,HSO4  +  2H2SO4  +  2HNO;! . 

Persulphates  act  as  strong  oxidizing  agents.  Salts  of  Ag' ,  Mn"  ,  Co"  ,  Ni"  and 
Pb"  ore  oxidized  in  the  presence  of  alkalies  to  the  peroxides  of  these  metals. 

If  ammonia  and  a  little  silver  nitrate  are  added  to  a  strong  solution  of  ammonium 
persulphate,  nitrogen  is  rapidly  evolved  and  the  solution  becomes  heated  to  boiling. 
The  silver  peroxide  first  formed  oxidizes  the  ammonia  with  liberation  of  nitrogen. 
(Z.  rtujs.  Ch.,  87,  255,  1901.) 

Fe"  and  Ce'"  are  oxidized  to  Fe'"  and  Ce""  salts.  KI  is  rapidly  oxidized; 
K4Fe(CN)6  becomes  K3Fe(CN)6  ;  Alcohol  is  slowly  oxidized  to  aldehyde,  rapidly 
on  warming;  organic  dyes  are  slowly  bleached. 

4.  Caio's  Acid. — By  adding  a  solid  persulphate  to  concentrated  sulphuric  acid 
at  0°,  a  solution  is  obtained  possessing  strong  oxidizing  properties.     It  may  also 
be  obtained  by  adding  30  per  cent  hydrogen  peroxide  (perhydrate)  to  concentrated 
sulphuric  acid: 

H2SO4  +  H2O2  =  H2O  +  H2SO5 . 

The  acid  HgSO5  is  known  as  monopersulphuric  acid  [Z.  angew.,  1898,  845:  Ber., 
34,  853  (1901);  41,  1839  (1909)]. 

5.  Detection. — Persulphates  are  tested  for  by  their  oxidizing  properties  and 
formation  of  the  peroxides  of  some  metals.     They  are   distinguished  from  hydro- 
gen peroxide  by  the  fact  that  persulphates  do  not  decolorize  potassium  perman- 
ganates and  do  not  produce  a  yellow  color  with  titanium  sulphate. 

§268.  Chlorine.     Cl  =  35.46.     Valence  one,  three,  four,  five,   and  seven. 

1.  Properties. — Molecular  weight,  70.92.  Vapor  density,  35.8.  The  molecule 
contains  two  atoms,  C12  .  Under  ordinary  air  pressure  it  liquefies  at  —33.6°  and 
solidifies  at  — 102°  (Olszewski,  M.,  1884,  5,  127).  Under  pressure  of  six  atmos- 
pheres it  liquefies  at  0°.  It  is  a  greenish-yellow,  suffocating1  gas,  not  com- 
bustible in  oxygen,  burns  in  hydrogen  (in  sunlight  combines  explosively), 
forming  HC1  .  On  cooling  an  aqueous  solution  of  the  gas  to  0°,  crystals  of 
C12.10H2O  separate  out  (Faraday,  Quart.  Jour,  of  Sci.,  1823,  15,  71).  Chlorine 
when  passed  into  a  solution  of  KOH  produces,  if  cold,  KC1  and  KC1O  ,  if  hot, 
KC1  and  KC103:  2KOH  +  CL  —  KC1  +  KC10  +  H20;  GKOH  +  3C12  =  5KC1  + 
KC1O.,  +  3H2O  .  Passed  into  an  excess  of  NH4OH  ,  NH4C1  and  N  are  formed: 
8NH4OH  +  3C12  =  6NH4C1  +  N_  +  8H,O;  if  chlorine  be  in  excess  chloride  of 
nitrogen  is  formed:  NH4OH  -f  3C12  =  NC13  +  3HC1  +  H2O  .  The  NC13  is  one 
of  the  most  dangerous  explosives  known;  hence  chlorine  should  never  be  passed 
into  NH4OH  or  into  a  solution  of  ammonium  salts  without  extreme  caution. 
Chlorine  bleaches  litmus,  indigo  and  most  other  organic  coloring  matter. 

The  three  elements,  chlorine,  bromine  and  iodine,  resemble  each  other  in 
almost  all  their  properties,  reactions  and  combinations,  differing  (as  do  their 
atomic  weights,  35.45,  79.95,  126.85)  with  a  regular  progressive  variation;  so 
that  their  compounds  present  themselves  to  us  as  members  of  progressive 
series.  In  several  particulars  fluorine  (atomic  weight,  19.05)  corresponds  to  the 
first  member  of  this  series  (£13). 

Two  oxides  of  chlorine  have  been  isolated:  CLO  ,  hypochlorous  anhydride 
(§270),  and  C102  ,  chlorine  dioxide.  The  latter  is  made  by  the  addition  of 
H2SO4  to  KC1O3  at  0°.  It  is  a  yellowish-green  gas,  condensing  at  0°  to  a  red- 
brown  liquid.  At  — 59°  it  becomes  a  crystalline  solid,  resembling  K,Cr207  .  It 
may  be  preserved  in  the  dark,  but  becomes  explosive  in  the  sunlight. 


338  CHLORINE.  §208,  2. 

The  most  important  acids  containing  chlorine  are  discussed  under  the 
sections  following.     They  are: 
Hydrochloric  acid,  HC1 . 
Hypochlorous  acid,  HC10  . 
Chlorous  acid,  HC102 . 
Chloric  acid,  HOICK . 
Perchloric  acid,  HC104  . 

2.  Occurrence. — It  does  not  occur  free  in  nature,  but  its  salts  are  numerous, 
the  most  abundant  being  NaCl  . 

3.  Formation. — (a)  By  the  action  of  HC1  upon  higher  oxides  as  indi- 
cated in  §269,  6A.  The  usual  class-room  or  laboratory  method  is  illus- 
trated by  the  following  equations: 

Mn02  +  4HC1  =  MnCL  +  C12  -f  2H,O 

Mn02  +  2NaCl  +  3H,SO4  =  MnSO4  +  2NaBSO4  +  01,  +  2H20 

(b)  By  fusing  together  NH4N03  and  NH4C1  :  4NH4N03  +  2NH4C1  =  5N,  + 
CL  +  12H2O  .  (c)  By  ignition  of  dry  MgCL  in  the  air:  2MgCl«  +  02  =  :  MgO 
+  2C12  (Dewar,  J.  8oc.  Ind.,  1887,  6,  775).  (d)  Some  chlorides"  are  disso<  iated 
by  heat  alone:  2AuCl3  =  2Au  +  8C12  . 

4.  Preparation. — (a)    Weldon's  process:  Mn02   is   treated   with  HC1  ,   an-i  the 
MnCL,  formed  is  precipitated  as  Mn(OH),  by  adding  Ca(OH)2  .     The  Mn  OH)2 
is  warmed  by  steam,  and  air  is  blown  into  it,  oxidizing  it  again  to  MnO,  ,  and 
by  repeating  this  process  the  same  manganese  is  used  over  again.     See  ]  unge 
and  Prett  (Z.  angew.,  1893,  99)   for  modification  of  this  method,  using  E  NO,  . 
(&)  Deacon's  process:  HC1  ,  mixed  with  air,  is  passed  over  fire-bricks  moistened 
with   CuCl2   and  heated  to  about  440°.     The   heat  first  changes  the   CuOL   to 
CuCl ,  evolving  chlorine;  then  the  oxygen  of  the  air,  aided  by  the  HC1  ,  oxi- 
dizes the  CuCl  to  CuCl2  .     It  is  not  certain   that  the   explanation   is   correct. 
It  is  only  known  that  the  hydrochloric  acid  which  is  passed  into  the  appj  ratus 
comes  out  as  free  chlorine,  and  that  the  copper  chloride   (small  in  amount) 
does  not  need  renewing,     (c)  Electrolytic  process:   Chlorine  is  very  largely  pro  luced 
as  a  by-product  in  the  manufacture  of  caustic  soda  by  the  electrolysis  of  common 
salt.     A  number  of  processes  have  been  developed  in  some  of  which  the  fuse  I  salt 
is  electrolyzed  while  in  others  the  electric  current  is  passed  through  strong  brine. 

5.  Solubilities. — The  maximum  solubility  of  chlorine  in  water  is  at  10°. 
At  0°  one  volume  of  water  dissolves  1.5  volumes  of  chlorine;  at  10°  three 
volumes;  at  30°  1.8  volumes   (Eiegel  and  Walz,  J.,  1846,  72).     B-iling 
completely  removes  the  chlorine  from  water.     The  chlorine  acts  upo:i  the 
water  to   a   small  extent,   forming  hydrochloric   and   hypochlorous   acid : 

C12  H-  H2O  s=>  H  Cl+H  CIO 

The  reaction  is  reversible,  but  if  an  alkali  is  present  the  hydrogen  ions 
are  removed  and  all  of  the  chlorine  reacts  with  the  water: 

C12  +  H2O  +  2OH  =  Cl  +  CIO  +  2H2O 

Only  chlorides  and  hypochlorites  remain  in  solution.  On  acidifying 
.the  solution  the  chlorine  is  again  liberated. 


§268,  651.  CHLORINE.  339 

6.  Reactions.  A. — With  metals  and  their  compounds. — Chlorine  is  one 
of  the  most  powerful  oxidizing  agents  known,  becoming  always  a  chloride 
or  hydrochloric  acid.  All  metals  are  at  lacked  by  moist  chlorine,  forming 
chlo  ides,  many  of  them  combining  with  vivid  incandescence.  With  per- 
fectly dry  chlorine  many  of  the  metals  are  not  at  all  attacked.  Sn , 
Sb  ,  :nd  As  are  rapidly  attacked,  forming  liquid  chlorides  (Cowper,  J.  (7., 
188c,  43,  153;  Veley,  J.  C.,  1894,  65,  1).  In  the  presence  of  acids  the 
oxid  ition  of  the  metal  takes  place  to  the  same  degree  as  when  that  metallic 
compound  is  acted  upon  by  HC1  (§269,  6A);  a  chloride  is  formed  having 
the  same  metallic  valence  that  would  have  resulted  from  treating  the 
oxide  or  hydroxide  with  hydrochloric  acid,  e.  g.,  adding  HC1  to  Co203  makes 
CoCl ,  not  CoCl3 ,  hence  adding  chlorine  to  metallic  cobalt  makes  CoCL  and 
not  GoClo, .  In  alkaline  mixture  usually  the  highest  degree  of  oxidation 
possible  is  attained,  as  indicated  by  the  following: 

1.  Pb"  becomes  Pb02  and  a  chloride  in  alkaline  mixture.    With  PbCl2 ,  it 
is  cli.imed  that  the  unstable  PbCl4  is  formed  (Sobrero  and  Selmi,  A.  Ch., 
1850,  (3),  29,  162;  Ditte,  A.  Ch.,  1881,  (5),  22,  566). 

2.  Hg'  becomes  Hg"  in  acid  and  in  alkaline  mixture;  also  HC1  or  a 
chloride. 

3.  As"'  becomes  Asv  in  acid  and  in  alkaline  mixture.     Some  water  must 
be  present  or  the  reverse  action  takes  place,  forming  AsCl.5  (§269,  6A2). 

4-  Sb"'  becomes  Sbv  and  a  chloride  with  acids  and  alkalis. 

5.  Sn"  becomes  SnIV  and  a  chloride  with  acids  and  alkalis. 

6.  MoVI~n  becomes  MoVI  and  a  chloride  with  acids  and  alkalis. 

7.  Bi"'  becomes  Biv  and  a  chloride  with  alkalis  only. 

8.  Cn'  becomes  Cu"  and  a  chloride  with  alkalis  and  with  acids. 

9.  Cr'"  becomes  CrVI  and  a  chloride  in  alkaline  mixture  only. 

10.  Fe"  becomes  Fe'"  and  a  chloride  with  acids  and  alkalis,  but  with 
alkalis  it  is  also  further  oxidized  to  a  ferrate. 

11.  Co"  becomes  Co(OH)3  and  a  chloride  with  alkalis  only. 

12.  Ni"  becomes  Ni(OH)3  and  a  chloride  with  alkalis  only. 

IS.  Mn"  becomes  Mn02  and  a  chloride  with  alkalis  only.  See  Ditte,  /.  c.s 
for  formation  of  MnCl4  . 

B. — With  non-metals  and  their  compounds. 

1.  H2C204  in  acid  mixture:  H2C204  +  C12  ==  2C02  +  2HC1 ,  the  H2C204 
must  be  in  excess  and  hot  (Guya-rd,  Bl.,  1879,  (2),  31,  299);  in  alkaline 
mixture:  K2C204  +  4KOH  +  C12  =  2K2C03  +  2KC1  +  2H20  . 

HCN  becomes  CNC1  and  HC1  (Bischoff,  B.,  1872,  5,  80). 

HCNS  forms  NH3 ,  H2S04 ,  C02 ,  and  other  variable  products,  and  HC1 
(Liebig,  A.,  1844,  50,  337). 

H4Fc(CN)6  becomes  H3Fe(CN)0  and  HC1  ;  an  excess  of  Cl  finally  decom- 
poses the  H3Fe(CN)6 . 


340  CHLORINE.  §268,  6£2. 

2.  Chlorine  does  not  appear  to  have  any  oxidizing  action  upon  the 
oxides  or  acids  of  nitrogen. 

3.  Phosphorus  and  all  lower  oxidized  forms  become  ILP04  with  forma- 
tion of  HC1  . 

4-  Sulphur  and  all  its  lower  oxidized  forms  are  oxidized  to  H2S04  with 
formation  of  HC1  .  In  an  alkaline  solution  a  sulphate  and  a  chloride  are 
formed.  With  H2S  ,  S  is  first  deposited,  which  an  excess  of  Cl  oxidizes  to 
H0S04  .  A  sulphide  in  an  alkaline  mixture  is  at  once  oxidized  to  a  sul- 
phate without  apparent  intermediate  liberation  of  sulphur. 

5.  In  alkaline  mixture  chlorine  oxidizes  chlorites,  and  hypochlorites  to 
chlorates  with  formation  of  a  chloride  :  KC102  +  2KOH  +  CL  ==  KC103 
+  2KC1  +  H20  .  With  NaOH  a  hypochlorite  is  formed  if  cold,  if  hot  a 

chlorate  : 

2NaOH  +  C12  =  NaCIO  +  NaCl  +  H20 
GNaOH  +  3C12  =  NaC103  +  SNaCl  +  3H2O 


6.  Chlorine  does  not  oxidize  bromine  in  acid  mixture,  in  alkaline  mix- 
ture a  bromate  and  a  chloride  are  formed.     HBr  in  acid  solution  becomes 
free  bromine,  in  alkaline  mixture  a  bromute  ;  hydrochloric  acid  or  a  chloride 
being  formed. 

7.  Iodine  is  oxidized  to  HI03  in  acid  mixture,  forming  HC1  ;  in  an 
alkaline  mixture  a  periodate  and  a  chloride  are  formed.     From  hydriodic 
acid  or  iodides,  iodine  is  first  liberated,  followed  by  further  oxidation  as 
indicated  above:  2HI  +  C12  =  2HC1  +  I2  ;  I,  +  5C12  +  6H20  =  2HI03  + 
10HC1  ;  KI  -f  8KOH  +  4C12  =  KI04  +  8KC1  +  4H20  . 

By  comparing  the  oxidizing  action  of  Cl  with  that  of  Br  and  I,  the 
following  facts  will  be  observed,  and  should  be  carefully  considered.  The 
elements  chlorine,  bromine,  and  iodine  have  an  oxidizing  power  in  reverse 
order  of  their  atomic  weights,  chlorine  being  the  strongest.  That  is,  if  all 
three  have  the  same  oxidizing  effect,  the  chlorine  acts  with  the  greatest 
rapidity;  and  in  some  cases,  as  with  cuprous  salts,  the  chlorine  oxidizes 
while  the  iodine  does  not.  Their  hydracids  are  reducing  agents  graded 
in  the  reverse  order.  If  any  increase  of  bonds  takes  place  in  presence  of 
an  acid,  by  chlorine,  bromine  or  iodine,  the  same  increase  always  occurs  in 
presence  of  a  fixed  alkali.  But  the  oxidation  frequently  goes  further  in 
presence  of  a  fixed  alkali.  Thus,  with  chlorine  and  potassium  hydroxide 
we  form  Pb02,  Ni(OH)3  ,  Bi205  ,  Co(OH)3  ,  K2Fe04,  and  Mn02  ,  which 
cannot  be  formed  in  presence  of  an  acid. 

It  is  very  important  to  remember  that  those  oxides  which  are  formed  ~by 
chlorine,  in  presence  of  a  fixed  alkali,  but  not  in  presence  of  an  acid,  are  the 
only  ones  which  can  be  reduced  by  hydrochloric  acid.  And  further,  that  this 
reduction  proceeds  not  always  to  the  original  form,  never  proceeding  beyond 
that  number  of  bonds  capable  of  being  formed  in  presence  of  an  acid.  Thus, 


§269,  30.  HYDROCHLORIC  ACID.  341 

any  lead  salt,  with  potassium  hydroxide  and  chlorine,  forms  PbO, ,  and 
this  treated  with  hydrochloric  acid  again  forms  the  lead  salt,  PbCl,  .  And 
ferrous  chloride  with  potassium  hydroxide  and  chlorine  forms  K2Fe04 ,  in 
which  iron  is  a  true  hexad,  and  K2Fe04  with  hydrochloric  acid  forms,  not 
the  ferrous  chloride  with  which  we  began,  but  ferric  chloride,  for  it  could 
only  be  oxidized  to  that  point  in  presence  of  an  acid. 

The  above  is  true  for  bromine  and  iodine,  as  well  as  for  chlorine. 

7.  Ignition. — See  1. 

8.  Detection. — Free  chlorine  is  recognized  by  its  odor,  by  its  liberation 
of  iodine  from  potassium  iodide,  by  its  bleaching  action  upon  litmus, 
indigo,  etc.,  and  by  its  action  as  a  powerful  oxidizing  agent  (see  above). 

Chlorine  acts  on  metallic  mercury  in  the  cold,  producing  the  insoluble 
mercurous  chloride: 

2Hg  +  C12  =  2HgCl . 

As  hydrochloric  acid  does  not  act  upon  metallic  mercury  it  may  be 
separated  from  chlorine  by  this  reaction,  the  mixture  being  shaken  with 
mercury  until  the  chlorine  is  removed.  Silver  nitrate  precipitates  one- 
sixth  of  the  chlorine  as  chloride : 

3C12  +  6AgNO3  +  3H2O  =  5AgCl+AgClO3  +  6HNO3 . 

9.  Estimation. — (a)  It  is  added  to  a  solution  of  potassium  iodide  and  the  lib- 
erated iodine  determined  by  standard  sodium  thiosulphate.     (6)  It  is  converted 
into  a  chloride  by  reducing  agents,  and  estimated  by  the  usual  methods  (§269,  8). 

§269.  Hydrochloric  Acid.    HC1  =  36.468. 
H'Cl-',  H—  01. 

1.  Properties. — Vapor    density,    18.22.     At    ordinary    pressure    it  liquefies  at 
-82.9°  and  solidifies  at  -112.5°  (Weber,  Z.  Anorg.,  74,  297  ;  1912).     At  10°  under 
pressure  of  40  atmospheres  it  condenses  to  a  colorless  liquid  (Faraday,  TV.,  1845, 
155).     Critical    temperature,    52.3°;     critical    pressure,    86    atmospheres    (Dewar, 
C.  N.,  1885,  61,  27).     Dissociated  into  H  and  Cl  at  about  1500°,  but   combines 
again  upon  cooling  (Deville,  C.  r.,  1865,  60,  317).     It  is  a  colorless  gas,  having 
an   acrid,   irritating  odor.     Readily   absorbed   by   water.     The   chemically   pure 
concentrated  acid  has  usually  a  specific  gravity  of  1.20,  and  contains  39.11  per 
cent  HC1  (Lunge  and  Marchlewski,  Z.  angew.,  1891,  4,  133).     The  U.  S.  P.  acid 
has  a  specific  gravity  of  1.163  at  15°  and  contains  31.9  per  cent  HC1  .     A  concen- 
trated solution  of  HC1  gives  off  gaseous  HC1  faster  than   H2O  ;  a  dilute  solution 
gives  off  H2O  faster  than  HC1  ,  as  a  final  result  in  both  cases  an  acid  sp.  (jr.  1.1 
distils  unchanged  at  110°  and  contains  20.18  per  cent  HC1  (Bineau,  A.  Cli.,  1843, 
(3),  7,  257). 

2.  Occurrence. — Found  native  only  in  the  vicinity  of  volcanoes.     Found  as  a 
chloride  in  many  minerals,  sodium  chloride  being  the  most  abundant. 

3.  Formation. — (a)  All  chlorides  except  those  of  mercury  are  trans- 
posed by  H2S04  ;  silver  chloride  must  be  heated  nearly  to  the  boiling  point 
of  the  H2S04  before  the  action  begins.  Lead,  antimony  and  tin  chlorides 
are  slowly  transposed. 


342  HYDROCHLORIC  ACID.  §2f9,  3&. 

(ft)  By  the  action  of  sunlight  on  a  mixture  of  H  and  Cl ,  or  by  heatmg  the 
mixture  to  150°.  (c)  Platinum  black,  palladium,  charcoal,  and  some  other  sub- 
stances which  rapidly  absorb  gases  will  cause  the  union  of  the  hydrogc  a  and 
the  chlorine,  (d)  When  hydrogen  is  passed  over  the  heated  chlorides  of  the 
most  of  the  metals  of  the  first  four  groups,  the  metals  are  set  free  and  "tydro- 
chloric  acid  is  formed,  (e)  Slowly  formed  by  the  action  of  chlorine  upon 
water  in  the  sunlight;  rapidly  by  its  action  upon  reducing  acids  srch  as 
H2C204 ,  HH2P02 ,  H2S,  H2S03 ,  etc.:  HH2PO2  +  2C12  +  2H2O  =  H3P04+  4HC1 . 

Chlorides  may  be  made:  (a)  By  direct  union  of  the  elements,  riostly 
without  heat.  Whether  an  ous  or  ic  salt  is  formed  depends  upca  the 
amount  of  chlorine  used,  (b)  By  the  action  of  hydrochloric  acid  upon  the 
corresponding  oxides,  hydroxides,  carbonates,  or  sulphites.  The  solutions 
formed  may  be  evaporated  to  expel  excess  of  acid.  If  the  chloride?  thus 
formed  contain  water  of  crystallization  it  cannot  be  removed  by  heat  alone, 
for  part  of  the  acid  is  by  this  means  driven  off,  and  a  basic  salt  rei  mins. 
If  the  anhydrous  chloride  is  desired,  it  may  always  be  made  by  (a1,  and 
when  thus  formed  may  be  sublimed  without  decomposition,  (c)  Chi  >rides 
of  the  first  group  are  best  made  by  precipitation,  (d)  Metals  solu  )le  in 
hydrochloric  acid  evolve  hydrogen  and  form  chlorides.  In  these  cases 
ous,  and  not  ic,  salts  are  formed,  (e)  Many  chlorides  may  be  formed  by 
bringing  HgCl2  in  contact  with  the  hot  metal. 

4.  Preparation. — For  commercial  purposes,  made  by  treating  NaC;  with 
H2S04  and  distilling. 

5.  Solubilities. — Hydrochloric  acid   (gas)  is  very  soluble  in  waier  as 
stated  in  (1);  forming  in  its  solutions  of  various  strengths  the  hydro- 
chloric acid  of  commerce.     Its  combinations  with  metals,  forming  chlor- 
ides, are  for  the  most  part  soluble  in  water.     AgCl  and  HgCl  are  insoluble 
in  water.     PbCL  is  only  slightly  soluble  in  cold  water  (§57,  5c).     These 
three  chlorides  constitute  the  first  or  silver  group  of  metals,  and  ar^  pre- 
cipitated from  their  solutions  by  hydrochloric  acid  or  soluble  chlorides 
(§61).   The  following  chlorides  not  commonly  met  with  are  inso'uble: 
cuprous   chloride  CuCl,  aurous  chloride  AuCl,  thallous   chloride  TlCi  and 
platinous    chloride,    PtCL .      The    following    oxychlorides    are    insoluble : 
BiOCl,  SbOCl  and  Hg,Cl20  .     Solutions  of  lead  salts  are  not  precipitated 
by  mercuric  chloride;  green  chromic  chloride  is  incompletely  precipitated 
and  a  sulphuric  acid  solution  of  molybdenum  oxychloride  not  at  dl  by 
silver  nitrate.     The  chlorides  of  Sb'" ,  Sn"  ,  and  Bi  require  the  presence 
of  some  free  acid  to  keep  them  in   solution.      AsCl3 ,  PC13 ,   SbCl .   and 
SnCl4  are  liquids  at  ordinary  temperature.     The  first  two  are  decorr  posed 
by  water  liberating  HC1  :  AsCl3  +  3H20  =±  H3As03  +  3HC1.    A  sat -rated 
solution  of  bismuth  nitrate  is  precipitated  by  HC1  as  the  oxychloride  ( §7  3,  6/) 
which  is  readily  soluble  in  excess  of  HC1 .     Hydrochloric  acid  increases 
the  solubility  of  the  chlorides  of  Pb  ,  Hg  ,  Ag  ,  Sb  ,  Au  ,  Pt ,  Bi  and  Cu'  ; 
it   decreases   the   solubility   of   Cd ,    Cu" ,    Co,    Ni ,    Mn,    Th ,    Ba,    Sr, 


§269,  §A2.  HYDROCHLORIC  ACID.  343 

Ca  ,  Mg  ,  Au  ,  K  and  NH4 .  Chlorides  of  Th  ,  Ba  ,  Na  ,  K  and  NH4  are 
nearly  insoluble  in  strong  HC1  ( Ditto,  C.  r.,  1SS1,  92,  242;  A.  Ch.f  1881, 
(5),  22,  551;  Berthelot,  A.  Ch.f  1881,  (5),  23,  8(5).  Chlorides  of  Li,  Ca 
and  Sr  are  soluble  in  absolute  alcohol  or  amyl  alcohol. 

Silver  chloride  is  readily  soluble  in  ammonium  hydroxide  (separation 
from  lead  and  mercurous  chlorides)  (§59,  6a);  lead  chloride  is  soluble  in 
fixed  alkali  hydroxides  (§57,  6a). 

H(  1  dissolves  or  transposes  all  insoluble  oxalates,  carbonates,  hypophos- 
phitcs,  phosphates,  and  sulphites.  Sulphides  of  Fe",  Mn ,  and  Zn  are 
dissolved  readily;  those  of  Pb ,  Ag,  Sb ,  Sn ,  Bi ,  Cu ,  Cd ,  Co,  and  Ni  if 
the  rcid  be  concentrated;  As2S3  and  AsaSr,  are  insoluble  in  the  cold  con- 
centrated acid,  very  slowly  soluble  in  the  hot  concentrated  acid;  HgS  , 
red,  is  insoluble;  black,  very  slowly  soluble  in  the  hot  concentrated  acid. 
HgS04  is  only  partially  transposed  by  HC1  (§58,  6/),  BaS04  not  at  all. 
The  ;nsoluble  sulphates  of  Pb ,  Hg',  Sr ,  and  Ca  are  slowly  but  completely 
dissolved  by  the  hot  concentrated  acid.  Many  of  the  metallic  chlorides 
are  soluble  in  alcohol,  a  few  are  soluble  in  ether. 

6.  Reactions. — A. — With  metals  and  their  compounds. — Hydrochloric 
acid  acts  upon  the  following  metals,  forming  chlorides  with  evolution  of 
hydrogen :  Pb  (slowly  but  completely),  Sn  ,  Cu  (very  slowly),  Cd ,  Fe  ,  Cr , 
Al .  C!o ,  Ni ,  Mn ,  Zn ,  and  the  metals  of  the  fifth  and  sixth  groups : 
Ag ,  Hg ,  As  ,  Sb  ,  Au  ,  Pt ,  and  Bi  are  insoluble  in  HC1  (Ditte  and  Metzner, 
A.  Ch.,  1893,  (6),  29,  389). 

The  following  metallic  oxides  and  hydroxides  are  acted  upon  by  hydro- 
chloric acid,  forming  chlorides  of  the  metal  without  reduction,  water  be- 
ing the  only  by-product :  Pb"  ,  Ag ,  Hg ,  As'"  (only  with  very  concentrated 
acid),  Sb,  Sn,  Au'",  Pt ,  MoVI,  Bi'",  Cu,  Cd,  Fe ,  Al ,  Cr"',  Co",  Ni", 
Mn".  Zn  ,  Ba ,  Sr ,  Ca ,  Mg ,  K  ,  and  Na  .  The  ignited  oxides  unite  with 
HC1  more  slowly  than  when  freshly  precipitated  or  when  dried  at  100°. 
Ignited  Cr,03  is  insoluble  in  HC1  :  other  ignited  oxides,  as  Fe203 ,  A1203 , 
etc.,  require  very  long  continued  boiling  with  the  HC1  to  effect  solution. 

The  following  metallic  compounds  are  attacked  by  hydrochloric  acid 
with  reduction  of  the  metal  and  evolution  of  chlorine: 

1.  Pb"+n  becomes  PbCL  ;  no  action  with  a  chloride  in  presence  of  a 
three  per  cent  solution  of  acetic  acid,  while  bromine  is  completely  set 
free  from  a  bromide  by  Pb02  in  presence  of  three  per  cent  of  acetic  acid 
(detection  of  a  chloride  in  presence  of  a  bromide)  (Vortmann,  M .,  1882,  3, 
510;  B.,  1887,  15,  1106). 

2.  Asv  becomes  AsCl3 .     (The  presence  of  very  concentrated  HC1   is 
required;  Fresenius,  Z.,  1862,  1,  448;  Smith,  J.  Am.  Soc.,  1895,  17,  682 
and  735.) 


344  HYDROCHLORIC   ACID.  §265,  6 AS. 

3.  Biv  becomes  BiCl, . 

4-  Crvl  becomes  CrCl3 .     With  K2Cr207 ,  bromine  is  completely  liberated 
from  a  bromide  in  presence  of  4  cc.  of  H2S04  to  100  cc.  of  water.     The 
chlorine  of  a  chloride  is  not  liberated,  and  the  bromine  may  be  removed 
by  boiling.     Test  the  solution  for  a  chloride  (Dechan,  J.  C.,  1886,  49, 
682).     Dry  HC1  does  not  reduce  CrVI  but  combines  with  it  to  form  the 
volatile  Cr02Cl2 ,  chlorochromic  anhydride  (method  of  detecting  a  chloride 
in  the  presence  of  a  bromide). 

5-  With  the  exception  of  ferrates  the  salts  of  iron  are  not  reduced  by 
hydrochloric  acid. 

6.  Co"+n  becomes  CoCl2 . 

7.  Ni"+n  becomes  NiCl2 . 

8.  Mn"+n  becomes  MnCL  .     Mn02  with  small  amounts  of  dilute  H2S04 
(1-10)  may  be  used  to  detect  a  chloride  in  presence  of  an  iodide  or  bromide. 
Boiling  the  mixture  removes  the  iodine  first,  then  the  bromine;  while  the 
chlorine  is  not  set  free  until  considerable  H2S04  has  been  added  (Jones, 
C.  N.,  1883,  48,  296).     A  mixture  of  KHS04  and  KMn04  completely  liber- 
ates the  bromine  from  a  bromide  in  the  cold.     A  chloride  remains  unde- 
composed  until  warmed.     Aspirate  off  the  bromine,  warm  and  collect  the 
chlorine  (Berglund,  Z.,  1885,  24,  184). 

B. — With  non-metals  and  their  compounds. 

1.  No  reducing  action  with  H2C204 ,  H2C03 ,  HCN,  HCNS,  H4Fe(CN)6, 
and  H,Fe(CN)e . 

2.  HN02  forms  chiefly  NO  and  Cl .     HN03  forms  N02C1  and  Cl ,   or 
NOC1  and  Cl ,  or  merely  N02  and  Cl .     In  case  excess  of  HC1  is  used  the 
reaction  is:  2HN03  +  6HC1  =  2NO  +  3C12  +  4H20  (Koninck  and  Nihoul, 
Z.  anorg.,  1890,  477).     Dry  HC1  gas,  passed  into  a  cold  mixture  of  con- 
centrated H2S04  and  HN03 ,  reacts  according  to  the  following  equations: 
2HC1  +  2HN03  =  2H20  +  2N02  +  C12  (Lunge,  Z.  angew.,  1895,  4,  8, 
and  11). 

3.  No  reducing  action  with  H2S ,  H2S03 ,  or  H2S04 .     With  thiosulphates 
the  unstable  H2S203  is  liberated  which  decomposes  as  follows :  2Na2S203  -(- 
4HC1  =  4NaCl  +  S2  .+  2S02  +  2H20  .     Sulphates  of  Ag  and  Kg'  are 
completely  transposed  by  HC1 ,  those  of  Ba ,  Sr ,  and  Ca  not  at  all,  all 
others  partially  (Prescott,  C.  N.,  1877,  36,  179). 

4.  With  an  excess  of  HC1,  hypophosphites,  phosphites,  and  phosphates 
are  dissolved  or  transposed  without  reduction. 

5.  Hypochlorous  acid  forms  chlorine  and  water :  HC10  +  HC1  =  H20  -j- 
C12 .      Chloric  acid  forms  C102 ,   C120 ,   and   Cl  in  varying  proportions, 
but  with  HC1  in  excess  the  following  reaction  takes  place :  KC103  +  6HC1 
;=  KC1  +  3C12  +  3H20  (Koninck  and  Nihoul,  Z,  anorg.,  1890,  481). 


§269,  8c.  HYDROCHLORIC  ACID.  345 

6.  KBr03  is  decomposed  by  boiling  with  HC1 ,  the  bromine  being  set 
free:    2KBrO,  +  12HC1  =  2KC1  +  Br,  +  5C12  +  6H,0  (Kaemmerer, 
J.  pr.,  1862,  85,  452). 

7.  With  HI03  ,  ICl-j  and  Cl  are  formed,  no  action  in  dilute  solutions: 
HI03  +  5HC1  =  IC13  +  Cl,  +  3H20  (Ditte,  A.,  1870, 156,  336).    According 
to  Bugarsky  (Z.  anorg.,  1895,  10,  387)  KHI20G  with  dilute  H2S04  does  not 
liberate   chlorine   from   a   chloride   even   on   boiling   (separation  from   a 
bromide). 

7.  Ignition. — The  chlorides  of  metals  are,  generally,  more  volatile  than  the 
other  compounds  of  the  same  metals:  example,  ferric  chloride. 

Insoluble  chlorides  are  readily  transposed  by  fusion  with  sodium  carbonate: 
PbCL  +  Na,C03  =  PbO  +  2NaCl  +  CO2  .  If  the  carbonate  be  mixed  with 
charcoal,  or  if  the  fusion  is  done  on  a  piece  of  charcoal,  the  metal  is  also 
reduced:  2PbCL  +  2Na2CO,  +  C  =  2Pb  +  4NaCl  +  3CO2  . 

Heated  in  a  bead  of  microcosmic  salt,  previously  saturated  with  copper 
oxide  in  the  inner  blow-pipe  flame,  chlorides  impart  a  blue  color  to  the  outer 
flame,  due  to  copper  chloride. 

Dry  sodium  sulphate  at  150°  is  transposed  by  dry  HC1  (Colson,  C.  r.,  1897, 
124,  81).  Gaseous  HC1  transposes  potassium  and  sodium  sulphates  completely 
at  a  dull-red  heat.  With  the  sulphates  of  the  alkaline  earths  the  transposition 
is  nearly  complete  (Hensgen,  B.,  1876,  9,  1671).  The  silver  halides  heated  with 
bismuth  sulphide  on  charcoal  before  the  blow-pipe  give  distinguishing  colored 
incrustations:  Agl ,  bright  red;  Ag-Br  ,  deep  yellow;  AgCl ,  white  (Goldschmidt, 
C.  C.,  1876,  297). 

8.  Detection. — (a)  In  its   soluble   compounds,   when   not   in   mixtures 
with  bromides  and  iodides,  hydrochloric  acid  is  readily  detected  by  pre- 
cipitation with  solution  of  silver  nitrate,  as  a  white  curdy  precipitate, 
opalescence  if  only  a  trace  be  present,  turning  gray  on  exposure  to  the 
light. 

The  properties  of  the  precipitate  of  silver  chloride  are  given  in  §59,  5c 
and  6/.  It  is  of  analytical  interest  in  that  it  is  freely  soluble  in  ammonium 
hydroxide  (considerably  more  freely  than  the  bromide,  and  far  more  freely 
than  the  iodide  of  silver);  soluble  in  hot,  concentrated  solution  of  am- 
monium carbonate  (which  dissolves  traces  of  bromide,  and  no  iodide  of 
silver);  insoluble  in^iitric  acid,  temporarily  soluble  in  strong  hydrochloric 
acid,  precipitating  again  on  dilution.  It  should  be  observed,  that  it  is 
appreciably  soluble  in  solutions  of  chlorides. 

(b)  A  test  for  traces  of  free  hydrochloric  acid,  in  distinction  from  metallic 
chlorides,  is  made  by  heating  the  solution  with  MnO,  ,  without  adding  an 
acid,  and  distilling  into  a  solution  of  potassium  iodide  and  starch.     Larger 
proportions  of  HC1  are  more  frequently  separated  by  distilling  it  intact. 

(c)  Gaseous  hydrochloric  acid  (formed  by  adding  sulphuric  acid  to  dry 
chlorides,  3a)  is  readily  detected  by  the  white  fumes  formed  when  brought 
in  contact  with  ammonia  vapor.     Also  by  bringing  a  stirring  rod  moist- 

'd  with  silver  nitrate  in  contact  with  the  hydrochloric  acid  gas,     Con- 


346  HYDROCHLORIC  ACID.  §269,  8d. 

firm  by  proving  the  solubility  of  the  white  precipitate  in   ammonium 
hydroxide. 

(d)  The  reaction  with  chromic  anhydride  is  in  use  as  a  test  for  hydro- 
chloric acid,  more  especially  in  presence  of  bromides: 

(a)     2HC1  +  Cr03  =  Cr02Cl2  (chlorochromic  anhydride)  +  H2O 
(ft)     4NaCl  +  K2Cr207  +  3H2S04  = 

2Cr02Cla  +  2Na2S04  +  K2S04  +  3H20 

To  obtain  a  rapid  production  of  the  gas,  so  that  it  may  be  recognized 
by  its  color,  the  operation  may  be  made  as  follows:  Boil  a  mixture  of 
solid  potassium  dichromate  and  sulphuric  acid,  in  an  evaporating-dish 
until  bright  red,  and  then  add  the  substance  *  to  be  tested,  in  powder- 
obtained,  if  necessary,  by  evaporation  of  the  solution.  If  chlorides  are 
present,  the  chromium  dioxydichloride  rises  instantly  as  a  bright  brownish- 
red  gas.  The  distinction  from  bromine  requires,  however,  that  the  mate- 
rial, which  should  be  dry,  should  be  distilled,  by  means  of  a  tubulated 
flask  or  small  retort,  the  vapors  being  condensed  in  a  receiver,  and  neutral- 
ized with  an  alkali  (c  and -d).  The  chromate  formed  makes  a  yellow  solu- 
tion (bromine,  a  colorless  solution).  As  conclusive  evidence  of  chlorine, 
the  chromate  (acidified  with  acetic  acid),  with  lead  acetate,  forms  a  yellow 
precipitate  (bromide,  a  white  precipitate,  if  any) : 

(c)  Cr02Cl2  +  2H20  =:  H2Cr04  +  2HC1 

(d)  CrO2Cl2  +  4(NH4)OH  =  (NH4)2Cr04  +  2NH4C1  +  2H20 

(e)  To  detect  a  chloride  in  the  presence  of  a  cyanide  or  thiocyanate, 
add  an  excess  of  silver  nitrate,  filter  and  wash.     To  the  moist  precipitate 
add  a  few  drops  of  silver  nitrate  (§318,  24)  and  then  several  cubic  centi- 
meters of  concentrated  sulphuric  acid  and  boil  for  two  or  three  minutes. 
The  silver  cyanide  and  thiocyanate  are  completely  dissolved  with  decom- 
position, while  the  silver  chloride  is  not  changed  except  on  long  continued 
boiling.     The  student  should  confirm  by  tests  on  known  material. 

According  to  Borchers  (C.  N.,  1883,  47,  218),  to  detect  a  chloride  in 
the  presence  of  a  cyanide  or  a  thiocyanate  add  silver^iitrate,  filter,  wash, 
and  boil  the  precipitate  with  concentrated  nitric  acid  to  complete  oxida- 
tion of  the  cyanogen  compound.  See  Mann  (Z.,  1889,  28,  668)  for  detec- 
tion of  a  chloride  in  presence  of  an  alkali  thiocyanate  by  use  of  CuS04 
and  H2S  . 

(/)  If  a  solution  containing  iodides,  bromides,  and  chlorides  be  boiled 
with  Fe2(S04)3 ,  all  the  iodine  is  liberated  and  may  be  collected  in  a 
solution  of  EJ  and  estimated  with  standard  Na2S203 .  The  solution  should 

*  With  the  chlorides  of  mercury  no  brown  fumes  are  obtained  as  these  chlorides  are  not 
transposed  by  the  sulphuric  acid;  and  the  chlorides  of  lead,  silver,  antimony,  and  tin  are  so 
slowly  transposed  that  the  formation  of  the  chromium  dioxydichloride  may  escape  observation. 
Before  relying  upon  this  test  the  absence  of  the  above  named  metals  should  be  assured. 


§269,  8k.  HYDROCHLORIC  ACID.  341} 

be  cooled  to  about  60°  and  a  slight  excess  of  KMn04  added.  The  bromine 
is  all  liberated  and  may  be  collected  in  NH4OH  and  estimated  as  a  bromide 
after  reduction  with  S02 .  The  chloride  may  now  be  detected  in  the 
filtrate  and  may  be  estimated  by  one  of  the  usual  methods.  Aspiration 
aids  the  removal  of  the  iodine  and  bromine  (Weiss,  C.  C.,  1885,  634  and 
712;  Hart,  C.  N.,  1884,  50,  268). 

(g)  Villiers  and  Fayotte  (C.  r.,  1894,  118,  1152,  1204  and  1413)  detect 
a  chloride  in  presence  of  an  iodide  and  bromide  by  passing  the  liberated 
halogens  into  a  solution  of  aniline  in  acetic  acid  (400  cc.  of  a  saturated 
water  solution  of  aniline  to  100  cc.  of  glacial  acetic  acid)  use  3  to  5  cc. 
of  this  solution  for  each  test.  Iodine  gives  no  precipitate;  bromine  gives 
a  white  precipitate;  and  chlorine  a  black  precipitate.  If  the  bromide  be 
present  in  large  excess,  add  silver  nitrate,  digest  the  precipitate  with 
ammonium  hydroxide,  add  hydrogen  sulphide  and  test  the  filtrate  as  the 
original  solution.  Liberate  the  halogen  with  KMn04  and  HQS04  . 

(70  Deniges  (Bl,  1890,  (3),  4,  481;  1891,  (3),  5,  66)  uses  H2S04  and 
Fe"'  to  liberate  the  iodine,  and  K2Cr04  to  liberate  the  bromine;  then 
after  boiling  off  the  I  and  Br  he  adds  KMn04  to  liberate  the  chlorine. 
The  iodine  he  detects  with  starch  paper,  the  bromine  fumes  are  absorbed 
on  a  rod  moistened  with  KOH ,  which  then  gives  an  orange-yellow  color 
with  aniline.  The  chlorine  he  collects  as  the  bromine  and  obtains  a  violet 
color  with  aniline. 

(i)  Dechan  (J.  0.,  1886,  50,  682;  1887,  51,  690)  removes  iodine  of 
iodides  by  distilling  with  a  concentrated  solution  of  K2Cr207;  then  the 
bromine  of  bromides  by  adding  dilute  H2S04  and  again  distilling.  The 
chloride  is  precipitated  by  AgN03  after  dilution  and  addition  of  HN03 . 

(/)  Vortman  (If.,  1882,  3,  510;  Z.,  1886,  25,  172)  detects  chlorine  in 
presence  of  bromine  and  iodine  as  follows:  The  solution  containing  the 
halogens  combined  with  the  alkali  or  alkaline  earth  metals  is  heated  with 
acetic  acid  and  peroxide  of  lead  until  the  supernatant  liquid  is  colorless 
and  has  no  longer  the  slightest  odor  of  iodine  or  bromine;  in  this  way  the 
whole  of  the  bromine  and  part  of  the  iodine  are  driven  off,  the  remainder 
of  the  latter  remaining  as  ioclate  of  lead  along  with  the  excess  of  lead 
peroxide.  This  is  filtered  off,  the  precipitate  washed  with  boiling  water, 
and  the  chlorine  precipitated  from  the  filtrate  by  addition  of  silver  nitrate. 

(k)  The  halogens  may  also  be  very  readily  separated  by  means  of 
potassium  persulphate,  K2S208  (§318,  15).  In  dilute  acetic  acid  iodine 
is  liberated  while  bromides  and  chlorides  are  not  oxidized.  On  acidifying 
with  H..S04  bromine  is  liberated,  while  there  is  no  action  on  chlorides  if 
the  strength  of  the  sulphuric  acid  does  not  exceed  2N.  If  the  free  iodine 
and  bromine  have  been  removed  by  CS2  or  boiling,  the  chlorine  may  be 


348  HYPOCHLOROUS  ACID.  §269,   9. 

precipitated  by  means  of  AgN03 .     As  in  the  presence  of  chlorates  the 
iodine  is  oxidized  to  I03 ,  C103  must  be  absent.     If  present,  the  halogens 
must  be  precipitated  as  silver  salts  and  reduced  with  metallic  zinc. 
The  following  reactions  take  place: 

2K1  +  K2S2O8  =  2K2SO4  +  I2 . 

2KBr  +  K2S2O8  +  H2SO4  =  2K2SO4  +  Br2  +  H2SO4 . 

KC1  +  K2S2O8  +  H2SO4(1.5— 2N)  =  No  action. 

2Agl  +  2AgBr  +  2AgCl  +  3Zn  =  6Ag  +  Znl2  +  ZnBr2  +  ZnCl2 . 

9.  Estimation. — (a) — It  is  precipitated  by  AgNO3 ,  washed,  and  after  igni- 
tion, weighed  as  AgCl .  (6)  By  a  standard  solution  of  AgNO3 .  A  little  Na2HPO4  , 
or,  better,  K2Cr2O7 ,  is  added  to  the  chloride  to  show  the  end  of  the  reaction. 
When  enough  AgNO3  has  been  added  to  combine  with  the  chlorine  the  next  addi- 
tion gives  a  yellow  precipitate  with  the  phosphate,  or  a  red  with  the  chromate. 


§270.  Hypochlorous  acid.    HC10  =  52.468  . 
H'Cl'O-",  H  —  0  —  Cl . 

1.  Properties. — Hypochlorous  anhydride,  CLO  ,  is  a  reddish-yellow  gas,  con- 
densing- at  about  — 20°  to  a  blood-red  liquid,  which  boils  at  about  — 17°  (Pelouze, 
A.   Ch.,   1843,    (3),   7,   176).     Rise   of  temperature   causes   decomposition,   explo- 
sively, into  chlorine  and  oxygen  (Balard,  A.  Ch.,  1834,  57,  225).     Molecular  iccnjltt, 
86.9.     Vapor  density,  43.5  at  10°.     The  acid,  HC1O  ,  has  not  been  isolated.     Its 
aqueous  solution  smells  like  C12O  ,  decomposing  rapidly,  especially  in  the  sun- 
light, into  Cl  and  HC103  . 

2.  Occurrence. — Not  found  in  nature. 

3.  Formation. — (a)    By   adding  chlorine  to  HgO  in  the   presence   of  water: 
2HgO  +  2C12  -f-  H20  =  Hg2OCl2  +  2HC1O  (Carius,  A.,  1863,  126,  196).     (b)  By 
adding  five  per  cent  nitric   acid  to   calcium   hypochlorite  and  distilling  at   a 
low  temperature  (Koffer,  A.,  1875,  177,  314).     (c)   By  passing  chlorine  into  the 
sulphates  of  Mg ,  Zn  ,  Al ,  Cu  ,  Ca  or  Na:  Na2SO4  +  C12  -f-  H2O  =  NaHS04  + 
NaCl  +  HC1O  .     (d)  By  heating  a  mixture  of  KC103  and  HoCoO4  to  70°  (Calvert 
and  Davies,  A.  Ch.,  1859,  (3),  55,  485). 

4.  Preparation. — For  commercial   purposes,   as  a  bleaching  agent  and   as  a 
disinfectant;  used  as  calcium  hypochlorite  with  calcium  chloride,  chlorinated 
lime,  made  by  bringing  chlorine  in  contact  with  calcium  hydroxide,  without 

heating.     Lunge   and   Schoch    (B.,   1887,   20,   1474)    give   the  formula   Ca"^1 

to  chlorinated  lime.  See  also  Kraut  (A.,  1882,  214,  244).  Also  as  sodium 
hypochlorite,  made  by  treating1  sodium  hydroxide  with  chlorine  short  of  satu- 
ration in  the  cold:  2NaOH  +  C12  =  NaCIO  +  NaCl  +  H2O  .  The  sodium 
hypochlorite-and-chloride — mixed  as  formed  by  chlorine  in  solution  of  sodium 
hydroxide  or  sodium  carbonate,  or  by  double  decomposition  between  solution 
of  the  calcium  hypochlorite-and-chloride  and  solution  of  sodium  carbonate — is 
pharmacopceial,  under  the  name  of  solution  of  chlorinated  soda  (NaCl. NaCIO). 

5.  Solubilities. — Hypochlorites  are  all  soluble  in  water  and  are  decomposed 
by  heating. 

6.  Reactions. — The  hypochlorites  are  all  unstable.     They  are  decomposed  by 
nearly  all  acids,  including  CO2:  2Ca(ClO)2    +   2CO2  =  2CaC03   +   2C12    +   O2; 


§271,  9.  CHLOROUS   ACID.  349 

4NaClO  +  4HC1  =  4NaCl  +  2H2O  +  2C12  -f  O? .  They  are  very  powerful  oxidiz- 
ing agents,  acting  in  acid  solution  as  free  chlorine,  as  the  above  equations  indicate. 
Hypochlorites  act  as  chlorine  in  alkaline  mixture  ($268,6)  (Fresenius,  Z.  angew., 
1895,  501).  On  warming,  all  hypochlorites  when  in  solution  are  converted  into 
chlorides  and  chlorates: 

3NaC10  =  2NaCl  +  NaClO3 . 

In  the  presence  of  40  per  cent  or  more  of  caustic  potash,  potassium  hypochlorite 
decomposes  into  chloride  with  evolution  of  oxygen  (Winteler,  Z.  Gngew.,  33, 
(1902),  778). 

When  shaken  with  mercury,  hypochlorites  or  free  hypochlorous  acid  produce  a 
reddish  basic  mercuric  chloride,  insoluble  in  water,  soluble  in  HC1  (distinction  from 
free  chlorine,  which  produces  white  mercurous  chloride  insoluble  in  HC1). 

7.  Ignition.- A 1 1  hypochlorites  are  decomposed  by  heat:  2KC10  =  2KC1  +  O2  . 

8.  Detection.-  Although    silver   hypochlorite   is   soluble    in    water,   it    decom- 
poses   very    quickly,   so   1h;i1    on    adding-   silver   nitrate   to   sodium   hypochlorite 
the  final  reaction  is  as  follows:  SNaCIO  +  :*AgNO3  =  2AgCl   +  AgC103   -f- 
3NaN03  .     When  KC1O  is  shaken  with  Hg°  ,  yellowish-red  Hg.OCL  is  formed; 
the  other  potassium   salts  of  chlorine,   i.e.,   KC1 ,   KC1O,  ,  KC103    and   KC10t  , 
have  no  action  upon  Hg°  .     An  indigo  solution  is  decolored  by  hypochlorites, 
while  KMnO4  is  not  decolored.     If  arsenons  acid  be  present,  the  indigo  solution 
is  not  decolored  until  the  arsenons  acid  is  all  oxidized  to  arsenic  acid. 

9.  Estimation. — It  is  estimated  as  AgCl  after  reduction  with  Zn  and  H.,S04  . 
Rosenbaum    (Z.  anyew.,  1893,   80)    gives  a   method  for   estimating  the   various 
chlorine  compounds  in  chlorinated  lime. 


§271.  Chlorous  acid.     HC102  =  68.468. 
H'Cl'"0-"2 ,  H  —  0  —  Cl  =  0  . 

1.  Properties.— The  anhydride,  C1203  ,  has  not  been  isolated  and  the  free  acid 
is  known  only  in  solution,  and  this  generally  contains  some  HC1O3  .     It  has  an 
intense  yellow  color  and  is  very  unstable. 

2.  Occurrence. — Neither  the  acid  nor  its  salts  are  found  in  nature. 

3.  Formation — An  impure  chlorous  acid  is  said  to  be  formed  when  KC103  is 
treated  with  HN03  and  As,03  ,  C^H^O^  or  C6H8   (Millon,  A.  Ch.,  1843,  (3),  7, 
298;  Schiel.  A.,  1859,  109,  318;  Carius,  A.,  1866,  140,  317).     Chlorites  of  a  number 
of  metals  have  been  made  by  adding  the  bases  to  a  water  solution  of  the  acid; 
also  from  KC10,  by  transposition. 

4.  Preparation.— KC1O2  is  prepared  by  adding  an  aqueous  solution  of  C102  of 
known  strength  to  the  proper  quantity  of  KOH  ,  and  evaporating  in  a  vacuum 
The  crystals  of  KC1O3  which  are  formed  in  the  reaction  are  removed  and  the 
mother  liquor  is  crystallized  from  alcohol. 

5.  Solubilities. — All  chlorites  which  have  been  prepared  are  soluble  in  water, 
lead  and  silver  chlorites  sparingly  soluble. 

6.  Reactions. — Chlorouc  acid  or  potassium  chlorite  in  dilute  acid  solution  is 
a  powerful  oxidizing  agent,  acting  similar  to  chlorine. 

7.  Ignition. — Chlorites  when  heated  evolve  oxygen  and  leave  a  chloride,  or 
first  a  chloride  and  a  chlorate  (Brandau,  A.,  1869,  151,  340). 

8.  Detection. — A  concentrated  solution  of  a  chlorite  gives  a  white  precipitate 
with  silver  nitrate,  fairly  readily  soluble  in  more  water.     KMnO,  is  decolored, 
a  brown  precipitate  being  formed.     A  solution  of  indigo  is  decolored  even  in 
presence   of    arsenous    acid    (distinction    from    hypochlorous    acid).     Chlorites 
when  slightly  acidulated  give  a  transient  amethyst  tint  to  a  solution  of  ferrous 
sulphate. 

9.  Estimation. — By  reduction  to  chloride  and  estimation  as  such.    By  meas- 


350  CHLORINE  PEROXIDE— CHLORIC  ACID.  §272. 

uring  the  amount  of  ferrous  iron  oxidized  to  the  ferric  condition:  4FeSO4  +  HC1O2 
L-  2H2S04  =  2Fe2(SO4)3  +  HC1  +  2H2O  . 


§272,  Chlorine  Peroxide,    C102  =  67.46. 

ciivo-"2,  o==cl~°~~cl:=0  or  °  =  C1=:0*- 

Chlorine  peroxide,  C1O2  ,  at  ordinary  temperature,  is  a  dark  greenish-yellow 
gas.  In  concentrated  solution  it  has  very  much  the  odor  of  nitrous  acid. 
Cooled  in  a  mixture  of  ice  and  salt  it  condenses  to  a  bromine-red  liquid;  and 
in  a  mixture  of  solid  C02  and  ether  it  forms  a  mass  of  orange-yellow,  brittle 
crystals.  When  warmed  to  about  60°  it  explodes  with  violence.  In  direct 
sunlight  at  ordinary  temperature  it  decomposes  slowlj-  into  chlorine  and 
oxygen,  while  in  the  dark  it  is  quite  stable.  In  contact  with  many  substances, 
as  phosphorus,  sulphur,  sugar,  ether,  turpentine,  etc.,  it  explodes  at  ordinary 
temperature.  In  moist  condition  it  bleaches  blue  litmus-paper  without  pre- 
viously reddening  it. 

One  volume  of  water  absorbs  about  20  volumes  of  the  gas  at  4°  (Millon, 
A.  07*.,  1843,  (3),  7,  298).  The  solution  in  \vater  contains  HC102  and  HC103  . 

It  is  prepared  by  carefully  adding  KC103  to  cold  concentrated  H2S04;  the 
mixture  is  then  carefully  warmed  to  20°,  later  somewhat  higher.  The  gas  is  con- 
densed in  a  tube  cooled  by  a  mixture  of  ice  and  salt:  3KC1O3  +  2H2S04  = 
2KHSO4  +  KC1O4  +  H2O  +  2C1O2  (Millon,  1.  c.).  It  is  also  made  by  warming 
a  mixture  of  oxalic  acid  and  potassium  chlorate.  When  prepared  in  this  man- 
ner it  is  mixed  with  C02:  2KC1O,  +  2H2C2O4  =  K2C2O4  +  2H2O  +  2C1O2  + 
2CO2  (Calvert  and  Davies,  A.,  1859,  110,  344).  It  is  also  formed,  mixed  with 
chlorine,  when  KC1O3  is  warmed  with  HC1 .  HI  is  oxidized  to  I;  SO2  to  H2S04  . 
Indigo  is  bleached  even  in  presence  of  As2O3  . 

§273.  Chloric  acid.     HC103  =  84.468. 
H'ClvO-"3 ,  H  —  0  —  Cl  21  £ 

1.  Properties. — A  solution  of  chloric  acid  may  be  evaporated   in   a   vacuum 
until  its  specific  gravity  is  1.282  at  14°.     The  composition  is  then  HC103.7H2O  , 
containing  40.1  per  cent  HC1O3    (Kaemmerer,  Pogg.,  1869,   138,  390).     Farther 
attempts   at   concentration    on    heating   to    40°   result   in   evolution    of  chlorine 
and   oxygen,  forming    HC1O4  :  8HC1O3  =  4HC1O4  +  2H2O+3O2  +  2C12    (Serullas, 
A.  Ch.,  1830,  45,  270).     Its  solution  in  the   cold  is  odorless  and  colorless;    first 
reddening    and   then   bleaching   litmus.     It   is  a   strong   oxidizing   agent,   paper 
soaked  with  the  acid  takes  fire   on  drying.     The  anhydride,  C12O6  ,   has  not  been 
isolated. 

2.  Occurrence. — Does  not  occur  in  nature. 

3.  Formation. — The  free  acid  may  be  formed  by  adding  an  excess  of  H2SiF6 
to  a  hot   solution  of  KC1O3 ;    the  filtrate  is  evaporated  in  vacuo,  the  excess  of 
H2SiF6  volatilizes,  leaving  the  HC1O3 .     Many  chlorates  are  formed  by  treating 
the    metallic    hydroxides   with    the  free  acid.     Also  by  the  action  of  Ba(Clp3)2 
upon  the  sulphate  of  the  metal  whose  chlorate  is   required;    or   by   the   action 
of  the  chloride  of  the  chlorate  needed,  upon  a  solution  of  AgClO3 . 

*  Pebal,  A.,  1875,  177,  1. 


£273,  GBL  CHLORIC  ACID.  351 

4.  Preparation. — ijy  adum!*1  H,SO.i  in  molecular  proportions  1o  a  solution  of 
Ba(C103)2  .  Chlorates  of  the  fifth  and  sixth  group  metals  are  prepared  by 
passing1  chlorine  into  the  respective  hydroxides  dissolved  or  suspended  in  water. 
By  repeated  crystallization  the  chlorate  is  separated  from  the  chloride  which 
is  also  formed:  6KOH  +  3CL  =  5KC1  -f  KC103  +  3ELO  . 

5.  Solubilities. — All  chlorates  are  soluble  in  water,  the  chlorates  of 
Hg ,  Sn  ,  and  Bi  require  a  little  free  acid.  Mercurous  and  ferrous  chlorates 
are  very  unstable.  Potassium  chlorate  is  the  least  soluble  of  the  stable 
metallic  chlorates;  soluble  in  about  21  parts  water  at  10°  (Blarez,  C.  r.> 
1891,  112,  1213). 

(].  Reactions.  A. — With  metals  and  their  compounds. — Chloric  acid 
attacks  Mg  evolving  hydrogen  and  forming  a  chlorate  only.  With  Zn , 
Fe,  Sn,  and  Cu  some  chloride  is  also  formed.  With  Zn  and  H2S04  the 
reduction  to  chloride  is  complete,  and  with  sodium  amalgam  no  reduction 
whatever  (Thorpe,  J.  C.,  1873,  26,  541).  With  the  zinc-copper  couple  * 
the  reduction  to  a  chloride  is  rapid  and  complete.  The  hot  concentrated 
acid  attacks  all  metals.  With  oxides  or  hydroxides  the  acid  forms  chlor- 
ates provided  a  chlorate  of  that  metal  can  by  any  means  be  formed.  Free 
chloric  acid  is  a  strong  oxidizing  agent,  and  if  an  excess  of  the  reducing 
agent  is  used,  it  is  converted  into  hydrochloric  acid,  or  a  chloride.  With 
the  aid  of  heat  the  chloric  acid  splits  up,  forming  some  chlorine  and 
oxides  of  chlorine. 

Hg'  forms  Hg". 

As"'  forms  Asv. 

Sb'"  forms  Sbv. 

Sn"  forms  Sniv. 

Cu'  forms  Cu". 

Cr"'  forms  CrVI,  chromic  salts  are  readily  oxidized  to  chromic  acid  on 
boiling  with  KC103  and  HN03 . 

Fe"  forms  Fe'"  (a  distinction  from  perchloric  acid)  (Carnot,  C.  r.,  1896, 
122,  452). 

Mn"  forms  MnIV,  manganous  salts  are  rapidly  oxidized  to  Mn02  on  warm- 
ing with  KC103  and  HNO, . 

Salts  of  lead,  cobalt,  and  nickel  do  not  appear  to  be  oxidized  on  boiling 
with  KC103  and  HN03 . 

B. — With  non-metals  and  their  compounds. 

1.  H2C204  forms  C02  and  varying  proportions  of  Cl  and  HC1 .  Heat 
and  excess  of  oxalic  acid  favors  the  production  of  HC1  (Guyard,  Bl,  1879, 

*  Gladstone  and  Tribe's  copper-zinc  couple  is  prepared  by  treating  thin  zinc  foil  with  a  1  per 
cent  solution  of  copper  sulphate  until  the  zinc  is  covered  with  a  black  deposit  of  reduced  cop- 
per. When  washed  and  dried  it  is  ready  for  use. 


352  CHLORIC  ACID.  §273, 

(2),  31,  299).  All  oxalates  are  decomposed,  C02  and  a  chlorate  or  chloride 
of  the  metal  being  formed.  Carbonates  are  all  transposed. 

HCNS  forms  H2S04 ,  HCN ,  and  HC1 . 

H4Fe(CN)6  first  forms  H;iFe(CN)6  and  HC1 ;  a  great  excess  of  HC103 
decomposes  the  H,Fe(CN)6 . 

2.  HN02  forms  HN03  and  Cl .     Nitrites  are  transposed  and  oxidized, 
forming  chlorates  or  nitrates  of  the  metal. 

3.  PH3 ,  HH2P02 ,  and  H,P03  form  H,P04  and  HC1 .     Hypophosphites 
and  phosphites  are  transposed  and  then  oxidized,  H3P04  and  a  chlorate  or 
a  chloride  of  the  metal  being  produced. 

4.  SVI~n  forms  SVI  and  HC1  ;  that  is,  the  sulphur  of  all  compounds 
becomes  H2S04  with  formation  of  HC1 .     All  sulphides,  sulphites,  thio- 
sulphates,  etc.,  are  transposed,  forming  a  chlorate,  chloride,  or  sulphate 
of  the  metal. 

5.  HC1  in  excess  forms  only  Cl  and  H20  (§269,  6B5).     NaCl  warmed  with 
HC103  evolves  01 ,  leaving  only  NaC103  . 

6.  HBr  forms  Br  and  HC1 .     KBr  warmed  with  HC103  evolves  Br ,  leav- 
ing only  KC103 . 

7.  I  and  HI  form  HI03  and  HC1 .     Soluble  iodides  form  iodic  acid  or 
an  iodate. 

7.  Ignition. — All  chlorates  are  resolved  by  heat  into  chlorides  and 
oxygen:  2KC103  =  2KC1  +  302 .  Some  perchlorate  is  usually  formed  as 
an  intermediate  product:  2KC103  =  KC104  +  KC1  +  02  (Serullas,  A.  Ch., 
1830,  (2),  45,  270).  In  presence  of  various  metallic  oxides,  etc.,  the 
oxygen  is  separated  more  easily,  the  metallic  oxides  remaining  unchanged. 
With  manganese  dioxide,  the  oxygen  of  potassium  chlorate  is  obtained  at 
about  200°;  ferric  oxide,  platinum  black,  copper  oxide,  and  lead  dioxide 
may  be  used  (§242,  3).  If  chlorates  are  rapidly  ignited  some  chlorine  is 
given  off  (Spring  and  Prost,  PL,  1889,  (3),  1,  340).  When  triturated  or 
heated  with  combustible  substances,  charcoal,  organic  substances,  sulphu". 
sulphites,  cyanides,  thiosulphates,  hypophosphites,  reduced  iron,  etc.— 
chlorates  violently  explode,  owing  to  their  sudden  decomposition,  and  the 
simultaneous  oxidation  of  the  combustible  material.  This  explosion  is 
more  violent  than  with  corresponding  mixtures  of  nitrates. 

Alkali  chlorates  when  fused  with  an  alkali,  or  an  alkali  carbonate,  and 
a  free  metal  or  a  lower  oxide,  or  salt  of  the  metal,  generally  oxidizes  it  to 
a  higher  oxide,  or  to  a  salt  having  an  increased  number  of  bonds;  and 
the  chlorate  is  reduced  to  a  chloride — e.  g.,  Mnvl~n  becomes  MnVI .  That 
is,  any  compound  of  manganese  having  less  than  six  bonds  is  oxidized  to 
the  hexad  (a).  Cr'"  becomes  CrVI  (6).  Asv~n  becomes  Asv  (c).  Pblv~n 


CALIFORNIA   COLLEGE 
of  PHARMACY 

§274,  3.  PERCHLORIC  ACID.  353 

becomes  PbIV  (d).     Co'"-n  becomes  Co"'  (e).     CIV~n  becomes  C1V  (/).     Pv~n 
becomes  Pv  (g).     Iv~n  becomes  Iv  (h).     SVI~n  becomes  SVI  (t). 
(a)     3Mns04  +  18KOH  +  5KC10,  =  9K,Mn04  -f  5KC1  +  9H.O 
(6)     2CrCl3  +  lONaOH  +  NaC103  =  2Na2CrO4  +  TNaCl  +  5H2O 

(c)  3As4  +  36KOH  -f  10KC1O3  =  12K,As04  +  10XC1  +  18H2O 

(d)  3Pb3Oi  +  Na.COs  +  2NaC103  =  9Pb02  +  2NaCl  +  Na2CO3 

(e)  GCoCL  +  12KOH  +  KC1O3  =  3Co203  +  13KC1  +  6H20 
(0      3K,C4H4Oe  +  5KC103  =  5KC1  +  3K2COS  +  9C02  +  6H2O 

(g)     3Pb(H2P02)2  +  18KOH  +  5KC103  =  3Pb02  +  «K3F04  +  5KC1  +  15H2O 
(ft)     ZnI2  +  K,C03  +  2KC1O3  =  ZnO  +,  2KIO3  +  2KC1  +  C02 
(i)      3K,S50G  +  12K2C03  +  10KC1O3  =  15K2SO4  +  10X01  +  12CO2 

8.  Detection,  ^u)  Dry  chlorates  when  warmed  with  concentrated  sul- 
phuric acid,  detonate  evolving  yellow  fumes  :  3KC103  +  2H2S04  —  2KHS04 
+  KC104  +  2C102  +  H20  .  This  action  is  modified  by  reducing  agents; 
some  acting  rapidly,  increase  the  detonation;  others  acting  slowly,  lessen 
it.  (6)  HC103  ,  like  HN03  ,  decolors  indigo  solution  and  gives  colors  with 
brucine.  diphenylamine,  paratoluidine,  and  phenol  similar  to  those  formed 
by  HN03.  (c)  By  ignition  a  chloride  is  left:  2KC103  =  2KC1  +  302  . 
(d)  It  is  changed  to  a  chloride  by  nascent  hydrogen:  2KC103  +  ^^n  -|- 
7H2S04  ==  6ZnS04  +  K2S04  +  2HC1  +  6H20;  or  by  reducing  acids  or 
bases:  2KC103  +  H2S04  +  6HJ30,  =  K2S04  +  6H2S04  +  2HC1  .  The 
resulting  HC1  is  then  identified  in  the  usual  manner.  Chlorides,  if  origin- 
ally present,  should  first  be  removed  by  silver  nitrate. 

9.  Estimation.  —  (a)  Reduction  to  a  chloride  and  estimation  as  such.   (&)  Addi- 
tion   of   HC1   and   XI   and   estimation   of   the   liberated   iodine   with   standard 


§274.  Perchloric  acid.     HC104  =  100.468. 

=  0 

H'ClVII0-"4  .  H  —  0  —  Cl  =  0 

=  0 

1.  Properties.—  Specific  gravity,  1.782  at  15°.     The  anhydrous  HC104  is  a  color- 
less oily  liquid,  volatile  but  cannot  be  distilled  without  partial  decomposition, 
often  with  explosive  violence.     Only  its  solution  in  water  can  be  safely  handled. 
Paper,  charcoal,  ether,  phosphorus,  and  many  other  substances  when  brought 
in  contact  with  the  anhydrous  acid  take  fire.     The  dilute  acid  is  very  stable,  not 
being-  easily  reduced  (Berthelot,  A.  Ch.,  1882,  (5),  27,  214).     It  does  not  bleach, 
but  merely  reddens  blue  litmus  paper. 

2.  Occurrence.  —  Not  found  in  nature. 

3.  Formation.—  (a)    By    electrolysis    of    a    solution    of    Cl    or    HC1    in    water 
(Riche,  C.  r.,   1858,  46,   348).     (6)   KC1O4   is  formed   by  electrolysis  of  KC1O3  , 
using  platinum  electrodes   (Lidoff  and  Tichomiroff,  J.   (7.,   1883,  44,   149).     (c) 
KC103   is  heated  with   an   excess   of  H,SiF6  ,  after  cooling  and   filtering,   the 
filtrate  is  carefully   distilled    (Roscoe,  J.   C.,   1863,   16,   82;   A.,    1862,    121,   346) 
(d)  By  treating  the  sulphate  of  the  metal,  the  perchlorate  of  which  is  desired, 


354  BROMINE.  §274,  4, 

with  Ba(ClO4)2  in  molecular  proportions,  (e)  By  treating  the  chloride  of  the 
metal,  the  perchlorate  of  which  is  desired,  with  AgC104  in  molecular  propor- 
tions. 

4.  Preparation. — KC104  is  made  by  carefully  heating  KC103   until  no  more 
oxygen  is  evolved:  2KC103  =  KC1  -f  KC1O4  +  O2  (7).     The  residue  is  dissolved 
in  water  and  upon  cooling1  crystals  of  KC1O4   separate.     The  free  acid,  nearly- 
pure,  is  obtained  by  cautiously  distilling  KC104  with  concentrated  H2S04  . 

5.  Solubilities. — All  of  the  perchlorates  of  the  ordinary  metals  are  soluble 
in   water,    and   all   are   deliquescent    except    NHtClO4     KC1O4  ,    Pb(C104)2    and 
HgClO4    (Serullas,  A.  Oh.,  1831,  46,  362).     Potassium  perchlorate  is  soluble  in 
142.9  parts  of  water  at  0°,  in  52.5  parts  at  25°,  and  in  5  parts  at  100°    (Muir, 
C.  N.,  1876,  33,  15).     KC1O4  is  insoluble  in  alcohol   (distinction  from  NaC104) 
(Schloessing,  A.  Ch.,  1877,  (5),  11,  561). 

6.  Reactions. — Iron  and  zinc  evolve  hydrogen  when  treated  with  perchloric 
acid.     The    acid    reacts    with    the   hydroxides    of    many    metals   to    form    per- 
chlorates.    It  is  not  reduced  by  HCl ,  HNO3  ,  H2S  or  SOo  .     Iodine  is  oxidized 
to  HIO4  with  liberation  of  chlorine:  I2  +  2HC1O4  =  2HI04  +  C12  .     A  solution 
of  indigo  is  not  decolored  by  HC104  even  after  the  addition  of  HCl  (distinction 
from  all   other  oxyacids  of  chlorine).     It   is  not   reduced   by   the   zinc-copper 
couple  (distinction  from  chlorate).     Sodium  perchlorate,  NaClO4  ,  is  used  as  a 
reagent  to  precipitate  potassiuir  salts. 

7.  Ignition. — Perchlorates  strongly  ignited  evolve  oxygen  and  leave  a  chloride 
(§242,  3). 

8.  Detection. — In  presence  of  a  hypochlorite,  chlorite,  chlorate  and  chloride 
boil  thoroughly  with  HCl;  the  first  three  are  decomposed,  leaving  chloride  and 
perchlorate.     Remove  the  chloride  with  AgN03  and  fuse  the  evaporated  filtrate 
with  Na2C03  .     Dissolve  the  fused  mass  in  water  and  test  for  a  chloride;  its 
presence  indicates  the  previous  presence  of  a  perchlorate.     Perchlorates  may  also 
be  separated  from  the  other  chlorine  acids  by  passing   SO2  gas,  which  reduces 
all  the  chlorine  acids  excepting  perchloric  acid.     Blattner  &  Brasseur,  Ch.  Z.,  24, 
793. 

9.  Estimation. — (a)  After  being  changed  to  a  chloride  a,s  indicated  above,  it 
is  estimated  in  the  usual  manner.     (6)  It  is  fused  with  zinc  chloride   and  the 
amount  of  chlorine  liberated  measured  by  the  amount  of  iodine  set  free  from  a 
solution  of  potassium  iodide   (separation  from  chlorate,   chlorides  'and  nitrates). 
(c)  KC1O4  is  heated  to  200°  with  HPO3  and  KI  ;    the  iodine  liberated  showing 
the  amount  of  perchlorate  present  (Gooch  and  Kreider,  Am.  S.,  1894,  48,   33;    and 
1895,  49  ,  287). 


§275.  Bromine.     Br  =  79.92.     Valence  one  and  five. 

1.  Properties. — Molecular  weight,  159.8;  vapor  density,  80;  specific  gravity, 
3.18828  at  0°;  boiling  point,  58.7°.  At  —7.3°  it  becomes  a  brown  solid  (Burgess, 
Wash.  Acad.  of  Sc.,  1-18).  At  ordinary  temperatures  bromine  is  a  brown-red, 
intensely  caustic  liquid,  freely  evolving  brown  vapors,  corrosive  vapors  of  a  suf- 
focating chlorine-like  odor.  As  a  solid  it  is  still  darker  in  color.  It  reacts  with 
KOH  in  all  respects  similar  to  chlorine  (§268,  1).  Indigo,  litmus  and  most  other 
organic  coloring  matters  are  bleached.  A  solution  of  starch  is  colored  slightly 
yellow. 

Bromine  decomposes  hydrosulphuric  acid  with  separation  of  sulphur,  and  sub- 
sequent production  of  sulphuric  acid;  changes  ferrous  to  ferric  salts,  and  (in 
presence  of  water)  acts  as  a  strong  oxidizing  agent.  It  displaces  iodine  from  iodides, 
and  is  displaced  from  bromides  by  chlorine;  its  character  being  intermediate 
between  that  of  chlorine  and  that  of  iodine. 

No  oxides  of  bromine  have,  with  certainty,  been  isolated.  The  well-estab- 
lished acids  are:  Hydrobromic,  HBr;  hypobromous,  HBrO;  bromic,  HBrO3, 


§275,  6^4,  11.  BROMINE.  355 

2.  Occurrence. — Not  found  free  in  nature.     As  a  bromide  in  sea  water,   mother 
liquor  from  salt  wells,  mineral  springs,  and  in  a  few  minerals. 

3.  Formation — (a)  Hydrobromic  acid  or  any  soluble  bromide  is  wanned  with 
MnO2   and   H,S04  .     (b)   Any   soluble   bromide   is   treated  with   chlorine   water 
and  the  solution  warmtd. 

4.  Preparation. — The    bromine    of    commerce    is    obtained    chiefly    from    the 
mother  liquor  of  the  salt  works:  (a)  By  treating-  with  MnO.,  and  H',S04:  MgBr,, 
+  Mn02  '+  2H2S04  =  MgSp4   +  MnS04   +   Br3   +   2H20  .     (b)   By  leading-  a 
current  of  steam  and   chlorine  into  the  bottom   of   a  vessel  filled  with   coke, 
into  which  a  stream  of  the  mother  liquor  flows  from  above:  MgBr2   +   CL  = 
MgCl2   -f  Br2  .     (c)   By  adding-  to  the  mother  liquor  a  mixture  of  Mg(OH)2  , 
suspended  in  water  and  saturated  with  chlorine,  rendering  acid  and  distilling- 
in  a  current  of  steam:  Mg(C103)2   -f  GMgBr,   +  12HC1  =  7MgCl2   +  6H20   + 
6Br2  .     ((/)  By  electrolysis  of  the  mother  liquor  at  a  low  temperature  and  then 
distilling-  in  a  current  of  steam. 

Commercial  bromine  is  freed  from  chlorine  by  adding-  KBr  and  distilling-.     If 
iodine  be  present  it  is  first  removed  as  Cul . 

5.  Solubilities. — Bromine  dissolves  in  30  parts  of  water  at  15°,  forming  an 
orange-yellow  solution  (Dancer,  J.  C.,  1862,  15,  477).     Its  water  solution  is  not 
permanent,   but    slowly    decomposes:    2Br2  +  2H2O  =  4HBr  -f  O2 .      Much    more 
soluble  in  HC1 ,  HBr ',   Kbr ,  BaCl2 ,  SrCl2 ,   and  in   many  other  salts   than  in 
water.     Soluble  in  carbon  disulphide,   chloroform,   ether  and  alcohol.     Readily 
removed  from  its  solution  in  water  by  shaking  with  carbon  disulphide  or  chloro- 
form, imparting  a  brown  color  to  the  solvent. 

6.  Reactions.  A. — With  metals  and  their  compounds. — Bromine  unites 
directly  with  gold,  platinum,  and  all  ordinary  metals  to  form  bromides. 
It  combines  with  metallic  mercury  forming  the  insoluble  mercurous 
bromide.  Silver  salts  are  precipitated,  yellow-white,  as  bromide  and 
bromate  :  6AgN03  +  3Br2  +  3H20  ==  5AgBr  +  AgBr03  +  6HN03 .  In  the 
following  metallic  compounds  the  valence  of  the  metal  is  changed;  the 
bromine  being  reduced  to  HBr  or,  if  in  alkaline  mixture,  to  a  bromide.  The 
reaction  is  less  violent  than  with  chlorine. 

1.  Pb"  becomes  Pb02  in  alkaline  mixture  only. 

2.  Hg'  becomes  Hg"  in  acid  and  in  alkaline  mixture. 

3.  As'"  becomes  Asv  in  acid  and  in  alkaline  mixture.      With  AsH3  and 
a  solution  of  bromine  in  water  H3As03  is  first  formed,  and  if  the  bromine 
be  in  excess  the  final  products  are  H3As04  and  HBr  . 

4.  Sb'"  becomes  Sbv  in  acid  and  in  alkaline  mixture. 

5.  Sn"  becomes  SnIV  in  acid  and  in  alkaline  mixture. 

6.  Bi"'  becomes  Bi205  in  alkaline  mixture  only. 

7.  Cu'  becomes  Cu"  in  acid  and  alkaline  mixture. 
S.  Cr'"  becomes  CrVI  in  alkaline  mixture  only. 

9.  Fe"  becomes  Fe'"  in  acid  mixture;  in  alkaline  mixture  the  iron  is 
further  oxidized  to  a  ferrate,  HBr  or  a  bromide  being  formed. 

10.  Co"  becomes  Co"'  in  alkaline  mixture  only. 

11.  Ni"  becomes  Ni'"  in  alkaline  mixture  only  (Kilpius,  J.  C.,  1876, 
29,  742). 


356  BROMINE.  §275, 6A,  12. 

12.  MnIV~n  becomes  MnIV  in  alkaline  mixture  only. 
B. — With  non-metals  and  their  compounds. 

1.  H2C204  becomes  a  carbonate  and  a  bromide  in  alkaline  mixture.     An 
excess  of  hot  saturated  oxalic  solution  changes  Br  to  HBr . 

HCNS  forms,  among  other  products,  H2S04  and  a  bromide  in  acid  mix- 
ture, and  a  sulphate  and  a  bromide  in  alkaline  mixture. 

H4Fe(CN)6  in  acid  mixture  forms  H3Fe(CN)6  and  HBr ,  in  alkaline  mix- 
ture a  ferricyanide  and  a  bromide  (Wagner,  J.  C.,  1876,  29,  741). 

2.  HN02  becomes  HN03  and  HBr  if  dilute  and  cold. 

3.  PH3 ,  HH2P02  and  H3P03  become  H3P04  and  HBr  with  acids,  and  a 
phosphate  and  a  bromide  in  alkaline  mixture.     P  and  Br  unite  to  form 
PBr3  or  PBr5 ,  depending  upon  relative  amounts  of  the  elements  present. 
The  phosphorus  bromides  are  decomposed  by  water,  forming  HBr  and 
the  corresponding  acids  of  phosphorus. 

4.  S°,  H2S ,  H2S03 ,  H2S203 ,  SVI~n  becomes  H2S04  and  HBr  with  acids, 
a  sulphate  and  a  bromide  in  alkaline  mixture. 

5.  Br  does  not  act  as  an  oxidizing  agent  upon  the  compounds  of  chlorine, 
but  may,  at  low  temperatures,  combine  with  chlorine  to  form  a  chlorine 
bromide,  BrCl  (Bornemann,  A.,  1877,  189,  183). 

6.  In  alkaline  mixture  hypobromites  by  boiling  are  oxidized  to  bromates 
with  formation  of  a  bromide. 

7.  Iodine  becomes  an  iodate  and  a  bromide  in  alkaline  mixture;  the 
elements  may  combine  to  form  the  unstable  bromiodide,  IBr  (Bornemann, 
/.  c.).     HI  and  iodides  form  I  and  HBr ,  but  in  alkaline  mixture  an  iodate 
and  a  bromide  are  produced. 

7.  Ignition. — Warming1  drives  off  all  the  bromine  from  its  solutions  in  water 
or  other  solvents.     Heat  favors  all  reactions  with  bromine. 

8.  Detection. — Bromine  is  usually  detected  by  shaking  its  solution  in 
water  with  CS2 ,  which  dissolves  it  with  a  reddish-yellow  color;  if  present 
in  large  quantities  the  color  is  brown  to  brownish  black.     In  this  case 
a  large  excess  of  CS2  must  be  used  or  a  very  small  portion  of  the  unknown 
taken,  in  order  that  the  solution  be  dilute  enough  for  the  reddish-yellow 
bromine    color    to   be    distinguished    from    the    violet    color    of  iodine. 
Ether  or  chloroform  may  be  used  instead  of  carbon  disulphide,  but  the 
solution  is  of  a  paler  yellow.     Starch  solution  gives  a  yellow  color  with 
bromine,  but  the  reaction  is  less  delicate  than  with  CS2 . 

9.  Estimation. — (a)  The  bromine  is  made  to  act  upon  KI  ,  and  the  iodine 
which  is  liberated  is  estimated  by  standard  solution  of  Na2S,O3  .  (ft)  It  is 
estimated  by  the  amount  of  As2O3  which  it  oxidizes  in  alkaline  solution,  (f)  It 
is  converted  into  HBr  by  HZS  or  H2S03  ,  and  then  precipitated  by 
and  weighed  as  AgBr . 


§276,  QA.  BYDROBROMIC  ACID.  357 

§276.     Hydrobromic  acid.    HBr  =  80.928. 
H'Br-',  H  — Br. 

1.  Properties. — Molecular  weight,  149.9.    Vapor  density,  39.1.    A  colorless  gas, 
condenses  to  a  liquid  at  — 09°  and  solidifies  at  — 73°  (Faraday,  A.,  1845,  56,  155). 
Its  aqueous  solution  is  colorless  and  is  not  decomposed  by  exposure  to  the 
air.     The  specific  gravity  of  the  saturated  solution  at  0°  is  1.78;  containing  82.02 
per  cent  HBr,   or  very   nearly   HBr.H20  .     If  a   saturated   solution   is   boiled, 
chiefly  HBr  is  given  off,  and  if  a  dilute  solution  is  boiled,  chiefly  H20  is  given 
off,  until  in  both  cases  the  remaining  liquid  contains  47.38  to  47.86  per  cent 
of  HBr  ,  its  sp.  (jr.  1.485,  its  boiling  point  constant  at  126°,  and  its  composition 
almost   exactly   HBr.5H20  ,   which  distils  over  unchanged.     Its  vapor   density 
of  14.1  agrees  with  the  calculated  vapor  density  of  HBr.5H20  . 

2.  Occurrence. — Not  found  free  in  nature,  in  combination  as  bromides  in  sea 
water  and  in  some  minerals. 

3.  Formation. — (a)    By    action    of    bromine    upon    phosphorus    immersed    in 
water,  the  amorphous  phosphorus  is  preferred:  P4  -+-  10Br2  +  16H2O  —  4H3P04 
-f-  20HBr  .     (b)  By  action  of  H3P04  or  H2S04  on  KBr  (Bertrand,  J.  C.,  1876,  29, 
877).     (c)  By  transposition  of  BaBr2  by  cold  dilute  H2SO4  added  in  molecular 
proportions,     (d)   By  passing  a  mixture  of  Br  and  H  over  platinum  sponge. 
(e)  By  action  of  Br  on  H3PO2  .     (f)  By  adding  Br  to  Na2S03  . 

Metallic  bromides  are  formed:  (1)  By  direct  union  of  the  elements,  but  in  a 
few  cases  heat  is  required  to  effect  the  combination.  (2)  By  action  of  HBr 
upon  the  metallic  oxides,  hydroxides  and  carbonates.  (3)  Many  bromides  are 
formed  by  action  of  HBr  on  the  free  metal,  ous  salts  and  not  ic  being  formed. 
(4)  Bromides  of  the  first  group  are  best  made  by  precipitation.  (5)  Bromides 
of  K  ,  Na  ,  Ba  ,  Sr  and  Ca  are  made  by  the  action  of  bromine  on  their  hydrox- 
ides and  subsequent  fusion: 

6KOH  -f  3Br3  =  KBr03  +  5KBr  +  3H20 

2KBr03  (ignited)  =  2KBr  +  30a 

4.  Preparation — (a)  H2S  is  added  to  a  solution  of  bromine  in  water  until 
the  yellow  color  disappears;  the  solution  is  then  distilled.     The  first  portion 
of  the  distillate  is  rejected  if  it  contains  H2S,  and  the  latter  portion  if  it  con- 
tains H2S04   (liecoura,  C.  r.,  1890,  110,  784).     (&)  H2S04  is  added  to  a  concen- 
trated  solution  of  KBr;  after  twenty-four  hours  the  greater  portion  of  the 
KHSO4    has    crystallized    out.     The    remaining    liquor    is    then    distilled.     The 
product  usually  contains  traces  of  H2SO4.     (c)   By  passing  bromine  into  hot 
paraffme  (Crismer,  B.,  1884,  17,  649). 

5.  Solubilities. — Silver  and  mercurous  bromide  are  insoluble  in  water, 
lead  bromide  is  sparingly  soluble ;  all  other  bromides  are  soluble.     Hydro- 
bromic acid  and  soluble  bromides  precipitate  solutions  of  the  metals  of 
the  first  group,  lead  salts  incompletely.     Lead  bromide  is  less  soluble  than 
the  corresponding  chloride.     The  presence  of  soluble  bromides  increases 
the  solubility  of  lead  bromide.     A  small  amount  of  hydrobromic  acid 
decreases  its  solubility,  but  a  larger  excess  increases  it  (Ditte,  C.  r.,  1881, 
92,  718). 

In  alcohol,  the  alkali  bromides  are  sparingly  or  slightly  soluble :  calcium 
bromide,  soluble;  mercuric  bromide,  soluble;  mercurous  bromide,  insolu- 
ble. Silver  bromide  is  soluble  in  NH4OH  . 

6.  Keactions. — A. — With  metals  and  their  compounds. — Hydrobromic 
acid  dissolves  many  metals  with  the  formation  of  bromides  and  evolution 
of  hydrogen,  e.  g.,  Pb ,  Sn ,  Fe ,  Al ,  Co ,  Ni ,  Zn ,  and  the  metals  of  the 


358  HYDROBROMIC  ACID.  §276,6.1,7. 

calcium  and  the  alkali  groups.  It  unites  with  salt  forming  oxides  and 
hydroxides  to  produce  bromides  without  change  of  valence:  PbO  +  2HBr 
=  PbBr2  +  H20  .  But  if  the  valence  of  the  metal  in  the  oxide  or 
hydroxide  is  such  that  no  corresponding  bromide  can  be  formed,  then 
reduction  takes  place  as  follows : 

1.  Pb"+n  becomes  PbBr2  and  Br  . 

2.  Asv  becomes  As'"  and  Br .     The  HBr  must  be  concentrated  and  in 
excess,  and  the  Asv  compound  merely  moistened  with  water:  H3As04  4- 
2HBr  =  H3As03  +  Br2  +  H20  .     In  presence  of  much  water  the  reverse 
action  takes  place :  H3As03  +  Br2  +  H20  =  H3As04  +  2HBr . 

3.  Sbv  becomes  Sb'"  and  Br  . 

4.  Biv  becomes  BiBr3  and  Br . 

5.  FeVI  becomes  Fe'"  and  not  Fe" ,  and  Br . 

6.  CrVI  becomes  CrBr3  and  Br  (a  separation  from  a  chloride  if  the  solu- 
tion be  dilute)  (Friedheim  and  Meyer,  Z.  anorg.,  1891,  1,  407).     KBr  is  not 
decomposed  by  a  boiling  concentrated   solution  of  K2Cr207   (separation 
from  KI)  (Dechan,  J.  C.,  1887,  51,  690). 

7.  Co"+n  becomes  CoBr2  and  Br  . 

8.  Ni"+n  becomes  NiBr2  and  Br  . 

9.  Mn"+n  becomes  MnBr2  and  Br  (§269,  8;  Jannasch  and  Aschoff,  Z. 
anorg.,  1891,  1,  144  and  245).     KMn04  liberates  all  the  bromine  from  KBr 
in  presence  of  CuS04  (a  separation  of  bromide  from  chloride  (Baubigny 
and  Rivals,  C.  r.,  1897,  124,  859  and  954). 

Silver  nitrate  solution  precipitates,  from  solutions  of  bromides,  silver 
bromide,  AgBr,  yellowish-white  in  the  light,  slowly  becoming  gray  to 
black.  The  precipitate  is  insoluble  in,  and  not  decomposed  by,  nitric  acid, 
soluble  in  concentrated  aqueous  ammonia,  nearly  insoluble  in  concentrated 
solution  of  ammonium  carbonate,  slightly  soluble  in  excess  of  alkali  bromides, 
soluble  in  solutions  of  alkali  cyanides  and  thiosulphates.  It  is  slowly  decom- 
posed by  chlorine  in  the  cold,  rapidly  when  heated  in  a  stream  of  chlorine. 

Solution  of  mercurous  nitrate  precipitates  mercurous  bromide,  HgBr, 
yellowish-white,  soluble  in  excess  of  alkali  bromides. 

Solutions  of  lead  salts  precipitate,  from  solutions  not  very  dilute,  lead 
bromide,  PbBr2 ,  white. 

B. — With  non-metals  and  their  compounds. 

1.  H,Fe(CN)(.  becomes  H4Fe(CN)6  and  Br .     The  HBr  must  be  in  excess 
and  concentrated,  also  the  ferricyanide  should  be  merely  moistened  with 
water,  as  in  the  presence  of  much  water  the  reverse  action  takes  place: 
2K4Fe(CN)6  +  Br2  ==  2K3Fc(CN)a  +  2KBr . 

2.  HN02 ,  in  dilute  solutions,  no  action  (distinction  from  HI)  (Gooch  and 
Ensign,  Am.  S.,  1890,  140,  145  and  283). 

HN03  becomes  NO  and  Br  . 


§276,  8.  HYDItOBROMIC  ACID.  359 

3.  Phosphorus  compounds  arc  not  reduced. 

4.  H2S04  becomes  S02  and  Br  .     Both  acids  must  be  concentrated  and 
hot,  otherwise  the  reverse  action  takes  place:   S02  +  Br2  -j-  2H20  =  H.,SO.t 
-f-  <?HBr  .     With  H2S04 ,  sp.  gr.  1.41,  no  bromine  is  set  free  even  wlii'ii 
solution  is  boiled  (Feit  and  Kubierschky,  J.  Pliarm.,  1891,  (5),  24,  159). 
The  bromine  of  bromides  is  all  liberated  when  warmed  to  70°  or  80°  with 
ammonium  persulphate  (separation  from  a  chloride)  (Engel,  C.  r.,  1894, 
118,  1263). 

5.  Chlorine  liberates  bromine  from  all  bromides,  even  from  fused  silver 
bromide  (Nihoul,  Z.  angew.,  1891,  441). 

HC103  becomes  HC1  and  Br .  If  the  HC103  be  concentrated  other  pro- 
ducts may  appear. 

6.  HBrO  liberates  Br  from  both  acids ;  the  same  with  HBr03 . 

7.  HI03  becomes  I  and  Br . 

8.  Hydrogen  peroxide  liberates  the  bromine  from  hydrobromic  acid  at 
100°  (a  distinction  and  separation  from  chloride).     The  bromine  can  best 
be  removed  by  aspiration  (Cavazzi,  Gazzetta,  1883,  13,  174). 

7.  Ignition. — Some  bromides  can  be  sublimed  undecomposed  in  presence  of 
air;  e.  </.,  AsBr,  ,  SbBr,  ,  HgBr  and  HgBr,  .  Some  can  be  sublimed  only  by 
exclusion  of  air  and  moisture;  c.  (/.,  AlBr3  and  NiBr2  .  Bromides  of  sodium  and 
potassium  are  not  changed  by  heat.  Silver  bromide  melts  undecomposed. 
Many  bromides,  however,  are  more  or  less  decomposed  when  ignited  in  pres- 
ence of  air  and  moisture:  CuBr2  becomes  CuBr  and  Br  . 

8.  Detection. — Bromides  are  usually  oxidized  to  free  bromine,  which  is 
detected  by  its  physical  properties  and  by  its  color  when  dissolved  in 
CS2  (£275,  5).  The  oxidizing  agent  used  to  liberate  the  bromine  varies 
according  to  the  conditions.  Chlorine  is  more  commonly  employed  and 
acts  when  cold  (6B5).  A  large  excess  of  chlorine  is  to  be  avoided,  as  it 
decolorizes  bromine  solutions  with  formation  of  a  chlorbromide..  Nitric 
acid  when  dilute  acts  slowly  unless  hot.  H2S04 ,  dilute,  fails  to  oxidize 
the  HBr  even  when  -hot;  but  when  concentrated  and  hot  is  sometimes 
preferred.  If  chlorine  be  used,  the  mixture  if  alkaline  must  first  be 
acidified;  otherwise  a  colorless  bromate  will  be  formed,  free  bromine  not 
being  a  visible  intermediate  step  in  the  oxidation :  KBr  +  GKOH  +  3C12 
=  KBr03  +  6KC1  +  3H,0  .  If  an  iodide  be  present:  (a)  In  absence  of  a 
chloride  precipitate  with  silver  nitrate,  and  digest  the  precipitate  with 
NH.OH ,  which  will  dissolve  the  AgBr  and  none  of  the  Agl .  The  filtrate 
may  be  treated  with  HJ3,  which  precipitates  the  silver  as  AgJS ,  leaving 
the  bromine  in  the  filtrate  as  NH4Br ,  which  may  be  detected  in  the  usual 
way.  (&)  To  the  acid  mixture  add  chlorine  water  and  carbon  disulphide, 
shake  and  continue  the  addition  of  the  chlorine  water  until  the  violet 
color  of  the  iodine  solution  disappears,  when  the  brown  color  due  to  the 
bromine  may  be  observed:  2KI  +  2KBr  +  7C12  +  6H20  =  2HIO,  +  Br, 


360  H7POBROMOUS  ACID—BROMIC  ACID.  §276,  9. 

-|-  4KC1  +  10HC1 .  (c)  To  the  solution  from  which  the  bases  have  been 
removed  add  a  cold  saturated  solution  of  potassium  chlorate  and  dilute 
sulphuric  acid  (one  of  acid  to  four  of  water);  warm  until  the  solution  is 
of  a  pale  straw  color,  or  colorless  if  only  iodides  are  present.  It  may  be 
necessary  to  add  more  of  the  solution  of  potassium  chlorate  to  complete 
the  oxidation  of  the  iodine.  Dilute  the  solution  with  water,  cool,  and 
shake  with  carbon  disulphide.  See  also  §269,  8. 

6KI  +  6KBr  +  2KC1O3  +  7H2S04  =  3I2  +  3Br2  +  7K2SO4  +  2HC1  +  GH20 

6la  +  xBr2  +  10KC103  +  5H2SO4  +  6H20  =  12HI03  +  xBr2  +  5K2SO4  +  10HC1 

9.  Estimation. — (a)  It  is  converted  into  AgBr  ,  and  after  gentle  ignition 
weighed  as  such.  (6)  The  bromide  is  oxidized  to  free  bromine,  which  is 
passed  into  a  solution  of  KI  and  the  liberated  iodine  titrated  with  standard 
Na2S203  .  (c)  The  bromide  is  oxidized  to  bromine,  which  is  passed  into  an 
alkaline  solution  of  arsenous  acid.  The  excess  of  the  arsenous  acid  is  titrated 
with  a  standard  solution  of  KMn04  . 


§277.     Hypobromous  acid.     HBrO  =  96.928. 
H'Br'O-",  H  —  0  —  Br. 

The  anhydride,  Br20  ,  has  not  been  isolated.  The  acid,  HBrO  ,  is  a  very 
unstable  yellow  liquid,  a  strong  oxidizing  and  bleaching  agent.  The  hypo- 
bromites  are  less  stable  than  the  corresponding  hypochlorites.  The  calcium 
and  the  alkali  group  hypobromites  may  be  prepared  by  adding  bromine  to  the 
respective  hydroxides  in  the  cold.  The  free  acid  is  obtained  by  the  action  of 
bromine  upon  mercuric  oxide:  2HgO  +  2Br2  +  H2O  =  Hg2OBr2  +  2HBrO; 
also  by  the  action  of  bromine  upon  silver  nitrate:  AgN03  -f  Br2  -f  H,O  = 
AgBr  +  HBrO  +  HNO3  (Dancer  and  Spiller,  C.  N.,  1860,  1,  38;  1862,  6,  249). 
The  free  acid  as  an  oxidizing  agent  reacts  in  many  cases  similar  to  free 
bromine.  With  HBr  free  Br  is  obtained  from  both  acids  (Schoenbein,  J.  pr., 
1863,  88,  475). 


§278.     Bromie  acid.    HBr03  =  128.928. 
H'BrvO-",,H—  0—  Br  = 

1.  Properties. — The  anhydride,  Br2O5  ,  has  not  been  isolated:  and  the  acid, 
HBrO3  ,  is  known  only  in  solution.     It  is  a  colorless  liquid,  smelling  like  bro- 
mine.    It  is  a  strong  oxidizing  agent.     The  solution  of  HBrO3  is  decomposed 
upon   boiling,  but  by  evaporating  in   a   vacuum   a   solution   containing   about 
50  per  cent  of  the  acid  may  be  obtained. 

2.  Occurrence.— Neither  the  acid  nor  its  salts  are  found  in  nature. 

3  Formation.— (ff)  By  the  electrolysis  of  HBr  (Riche,  C.  r.,  1858,  46,  348). 
(1))  By  the  decomposition  of  AgBr03  by  Br:  SAgBrO,  +  3Br,  +  3H,O  =  5AgBr 
-4-  6HBrO3  .  (c)  An  alkali  bromate  is  made  by  adding  bromine  to  a  solution 
of  chlorine  in  sodium  carbonate  (Kaemmerer,  J.  pr.,  1862,  85,  452). 

4.  Preparation.— Bromates  of  Ba ,  Sr  ,  Ca  ,  K  and  Na  are  made  byj.he  action 
of  bromine  upon  the  respective  hydroxides  at  100°:  6KOH  +  3Br2  ~-  5KBr  + 
KBrO,  +  3HoO  .  The  free  acid  is  prepared  by  adding  dilute  H,SO4  in  slight 
excess  to  Ba(BrO8)a;  the  slight  excess  of  H3S04  being  removed  by  the  cautious 
addition  of  £&(OB)i . 


§278, 8.  BROMIC  ACID.  361 

5.  Solubilities.— AgBr03    is   soluble   in    123   parts   of   water   at    24.5° 
(Noyes,  Z.  phys.  Ch.,  1890,  6,  246).     Ba(Br03),  is  soluble  in  124  parts  of 
•yater  at  ordinary  temperature  and  in  24  parts  at  100°    (Kammelsberg, 
Pogg.,  1841,  52,  81  and  86).     With  the  exception  of  some  basic  bromates, 
all  other  bromates  are  soluble  in  water. 

6.  Reactions. — A. — With  metals  and  their  compounds. — Bromic  acid  is 
a  powerful  oxidizing  agent,  acting  in  most  respects  like  free  bromine. 
It  is  usually  reduced  to  hydrobromic  acid,  sometimes  only  to  free  bromine : 

1.  Hg'  becomes  Hg"  and  a  bromide. 

2.  As"'  becomes  Asv  and  a  bromide. 

3.  Sb'"  becomes  Sbv  and  a  bromide. 
4-  Sn"  becomes  SnIV  and  a  bromide. 

5.  Cu'  becomes  Cu"  and  a  bromide. 

6.  Fe"  becomes  Fe'"  and  a  bromide. 

7.  Mn"  becomes  Mn02  and  bromine. 

8.  Cr'"  becomes  H2Cr04  and  bromine. 

Silver  nitrate  precipitates  in  solutions  not  very  dilute,  silver  bromate, 
AgBr03 ,  white,  sparingly  soluble  in  wrater,  soluble  in  ammonium  hydroxide, 
easily  soluble  by  nitric  acid,  its  color  and  solubility  in  ammonium  hydroxide 
differing  a  little  from  the  bromide  (§276,  5).  It  is  decomposed  by  hydro- 
chloric acid  with  evolution  of  bromine — a  distinction  from  bromides  and 
from  other  argentic  precipitates. 

5.— With  non-metals  and  their  compounds. 

1.  H2C204  becomes  CO.,  and  Br .     An  excess  of  hot  H2C204  changes  the 
Br  to  HBr  (Guyard,  BL,  1879,  (2),  31,  299). 

HCNS  becomes  H2S04 ,  HBr  and  other  products. 

H4Fe(CN)6  becomes  H3Fe(CN)6  and  HBr .  An  excess  of  HBr03  carries 
the  oxidation  farther. 

2.  HN02  reduces  HBr03 ,  forming  HN03  and  Br  . 

3.  PH3 ,  HH2P02  and  H3P03  become  H3P04  and  HBr . 

4.  S  and  S02  become  H2S04  and  HBr  . 
H2S  forms  first  S  then  H2S04  . 

5.  HC1  becomes  Cl  and  Br  . 

6.  HBr  forms  Br  from  both  acids. 

7.  HI  becomes  I  and  Br .     With  an  excess  of  HBr03  the  products  are 
HI03  and  Br  (Kaemmerer,  I.  c.,  Wittstein,  Z.,  1876,  15,  61). 

7.  Ignition. — All    bromates    are    decomposed    upon    heating.     KBr03 , 
NaBr03  and  Ca(Br03)2  evolve  oxygen  and  leave  the  bromides.     Co(Br03)2 , 
Zn(Br03)2  and  other  bromates  evolve  oxygen  and  bromine,  leaving  an  oxide. 

8.  Detection. — The  bromine  is  first  liberated  by  some  reducing  agent 
that  does  not  carry  the  reduction  to  the  formation  of  HBr,    H2C?04  is  $ 


362  IODINE.  §278, 9. 

very  suitable  agent  for  this  purpose,  since  it  does  not  change  Br  to  HBr 
except  when  hot  and  concentrated.  The  Br  is  detected  by  CS2  (§275,  8). 

Sulphuric  and  nitric  acids  liberate  bromic  acid  from  metallic  bromates, 
the  HBr03  remaining  for  some  time  intact,  and  the  solution  colorless.  The 
gradual  decomposition  of  the  HBr03  is  first  a  resolution  into  HBr  and  0, 
and  as  fast  as  HBr  is  formed  it  acts  with  HBr03 ,  so  as  to  liberate  the 
bromine  of  both  acids.  Now,  if  the  solution  contained  bromide  as  well  as 
bromate,,  an  abundance  of  free  bromine  is  obtained  immediately  upon  the 
addition  of  dilute  sulphuric  acid  in  the  cold.  Hence,  if  dilute  sulphuric 
acid  in  the  dilute  cold  solution  does  not  color  the  carbon  disulphide,  and 
if  the  addition  of  solution  of  pure  potassium  bromide  immediately  develops 
the  yellow  color,  while  it  is  found  that  no  other  oxidizing  agent  is  present, 
we  have  corroborative  evidence  of  the  presence  of  a  bromate.  And,  if  we 
treat  a  solution  known  to  contain  bromide  with  dilute  sulphuric  acid  and 
carbon  disulphide,  and  obtain  no  color,  we  have  conclusive  evidence  of  the 
absence  of  bromates.  Hydrochloric  acid  transposes  bromates  and  quickly 
decomposes  the  bromic  acid,  liberating  both  bromine  and  chlorine. 

A  mixture  of  bromate  and  iodate,  treated  with  hydrochloric  acid,  fur- 
nishes bromine  without  iodine,  coloring  carbon  disulphide  yellow. 

The  ignited  residue  of  bromates,  in  all  cases  if  the  ignition  be  done  with 
sodium  carbonate,  will  give  the  tests  for  bromides. 

9.  Estimation. — The  bromate  is  reduced  to  free  bromine  or  to  a  bromide  and 
determined  as  such. 

§279.  Iodine.    1=  126.92.     Usual  valence  one,  five  and  seven  (§12). 

1.  Properties.— Specific  gravity,  4.948  at  17°  (Gay-Lussac).  Melting  point, 
114.2°.  Boiling  point,  184.35°  at  760  mm.  pressure  (Ramsay  and  Young-,  J.  ('., 
1886,  49,  453).  At  ordinary  temperature  iodine  is  a  soft  gray-black  crystalline 
solid  with  a  metallic  lustre.  The  thin  crystals  have  a  brownish-red  appear- 
ance. Precipitated  iodine  is  a  brownish-black  powder.  It  vaporizes  very 
appreciably  at  ordinary  room  temperature  with  a  characteristic  odor,  and  may 
be  distilled  with  steam.  The  molecule  of  iodine  vapor  under  about  800°  is  I2; 
above  that  temperature  dissociation  takes  place,  until  at  1700°  it  is  complete 
and  the  molecule  consists  of  single  atoms  (Biltz  and  Meyer,  B.,  1889,  22,  725). 
The  vapor  of  iodine  unmixed  with  other  gases  is  deep  blue,  mixed  with  air 
or  other  gases  it  is  a  beautiful  violet.  It  is  sparingly  soluble  in  water  to  a 
brown  or  yellowish-brown  solution,  which  slowly  bleaches  litmus  paper.  It 
stains  the  skin  yellow-brown.  The  solution  gradually  decomposes  in  the  sun- 
light with  formation  of  HI.  It  reacts  similarly  to  bromine  and  chlorine,  but 
with  much  less  intensity.  The  free  element  combines  with  starch,*  forming 
a  compound  of  an  intense  blue  color.  This  colored  body  is  quite  stable  in  the 
cold;  decolors  upon  warming,  the  color  returning  upon  cooling.  The  reaction 
of  iodine  with  starch  constitutes  a  very  delicate  reaction  for  the  detection  of 
the  presence  of  iodine.  It  also  serves  as  an  indicator  in  the  volumetric  estima- 
tion of  iodine,  as  all  reducing  agents  destroy  the  color  by  taking  the  iodine 
into  combination.  Combined  iodine  does  not  react  with  starch. 

*  The  compound  formed  when  iodine  unites  with  starch  is  regarded  by  Bondonrneau  (BL, 
1877,  (2),  28,  452)  as  an  addition  compound  of  the  composition  (<J6H]op6)2l.  Mylius  (Der.,  20, 
tiSK).  gives  the  formula  [C24H4oO2oI]4.HI  and  considers  it  the  hydriodic  acid  compound  of  an 
iodine  addition  compound  of  starch.  This  view  is  supported  by  the  fact  that  in  the  absence  of 
KI  or  other  iodide  iodine  does  not  give  a  blue  color  with  starch.  (Sonnes,  Z.,  34,  409.) 


,  GA,  5.  IODINE.  363 

Colorless  solutions  are  formed  by  all  the  alkali  hydroxides  with  iodine;  the 
fixed  alkali  hydroxides  forming  iodides  and  iodates.  With  ammonia  in  water 
solution  it  dissolves  more  slowly,  becoming  colorless;  the  solution  contains  the 
most  of  the  iodine  as  ammonium  iodide,  and  deposits  a  dark-brown  powder, 
termed  "  iodide  of  nitrayrn"  very  easily  and  violently  explosive  when  dry. 
According  to  Chattavvay  (Am.,  I'.MX),  24,  138)  this  compound  has  the  composi- 
tion N,H3I3  . 

The  anhydride  of  iodic  acid,  I2OB  ,  is  the  only  stable  compound  of  iodine  and 
oxygen.  The  chief  acids  of  iodine  are:  Hydriodic  acid,  HI;  iodic  acid,  HIO3; 
periodic  acid,  HI04  . 

Hypoiodous  acid  is  said  to  be  formed  by  the  action  of  alcoholic  iodine  upon 
freshly  precipitated  mercuric  oxide  (Lippmann,  C.  r.,  18G6,  63,  908).  Lunge  and 
Schoche  (#.,  1882,  15,  1883)  prepared  iodide  of  lime  which  seemed  to  contain 
calcium  hypoiodite,  Ca(IO)2  . 

2.  Occurrence Found  free  in  some  mineral  waters  (Wanklyn,  C.  N.,  1886,  54, 

300).     As  iodides  and  iodates  in  sea  water   (Sonstadt,  C.  N.,  1872,  25,  196,  231 
and  241).     In  the  ashes  of  sea  plants.     In  small  quantities  in  several  minerals, 
especially  in  Chili  saltpeter  as  sodium  iodate. 

3.  Formation. — From  iodides  by  nearly  all  oxidizing  agents:  2KI  -+-  Br2  = 
2KBr  +  I,;  and  from  iodates  by  nearly  all  reducing  agents:  2HI03  +  5H2C2O4 
=  I2  +  IOCO,  +  6H20. 

4.  Preparation. —  (a)   The  ashes  of  the  sea  plants  are  digested  in  hot  water 
and  from  the  nitrate  most  of  the  salts  removed  by  evaporation  and  crystalliza- 
tion.    The  iodides  remain  in  the  mother  liquor  and  from  this  the  iodine  is 
obtained  by  treatment  with  Mn02  and  H2S04  .     (ft)  The  sodium  iodate  in  the 
mother  liquor  of  the  Chili  saltpeter  is  reduced  with  S02  ,  the  iodine  precipitated 
as  Cul  with  CuSO4  .     From  the  precipitate  the  iodine  is  recovered  by  distilla- 
tion with  MnO2   and  H2SO4  .     By  far_the  greatest  portion  of  the  iodine  and 
iodides  of  commerce  is  obtained  from  the  Chili  saltpeter  deposits. 

5.  Solubilities, — It  is  soluble  in  about  5500  parts  water  at  10°  to  12° 
(Wittstein,  J '.,  1857,  123),  differing  from  Cl  or  Br  in  that  it  forms  no 
hydrate.  It  is  much  more  soluble  in  water  containing  hydriodic  acid  or 
soluble  iodides.  From  a  concentrated  solution  in  KI  the  compound  KI3 
has  been  obtained.  Iodine  dissolves  in  very  many  organic  solvents  as 
alcohol,  ether,  chloroform,  glycerol,  benzol,  carbon  disulphide,  etc.  Car- 
bon disulphide  readily  removes  the  iodine  from  its  solution  or  suspension 
in  water;  with  small  amounts  of  iodine  imparting  to  the  carbon  disulphide 
a  beautiful  violet  color,  with  large  amounts  the  CS2  solution  is  almost 
black. 

G.  Reactions. — A. — With  metals  and  their  compounds. — It  unites  slowly 
by  the  aid  of  heat  with  Pb  and  Ag;  more  rapidly  with  Hg,  As,  Sb,  Sn, 
Bi ,  Cu  ,  Cd  ,  Al ,  Cr ,  Fe  ,  Co  ,  Ni ,  Mn  ,  Zn  ,  Ba  ,  Sr ,  Ca ,  Mg ,  K  and  Na  . 

In  oxidizing  metallic  compounds  the  iodine  invariably  becomes  HI  or 
an  iodide,  depending  upon  whether  the  mixture  be  acid  or  alkaline.  It 
may,  however,  with  certain  substances  act  as  a  reducing  agent,  becoming 
oxidized  to  iodate  or  periodate. 

1.  Hg'  becomes  Hg"  in  acid  and  in  alkaline  mixture. 

2.  As'"  becomes  Asv  in  presence  of  alkalis  only. 

3.  Sb'"  becomes  Sbv  in  presence  of  alkalis  only. 

4-  Sn"  becomes  SnIV  in  acid  or  in  alkaline  mixture. 
5.  Cr'"  becomes  Crvl  in  presence  of  alkalis  only. 


364  IODINE.  §279,  6A,  6. 

6.  Fe"  becomes  Fe'"  in  presence  of  alkalis  only. 

7.  Co"  becomes  Co'"  in  presence  of  alkalis  only. 

8.  Ni"  is  not  oxidized. 

9.  Mn"  becomes  Mnlv  in  presence  of  alkalis  only. 
B. — With  non-metals  and  their  compounds. 

1.  K4Fe(CN)6  is  oxidized,  forming  K3Fe(CN)6  and  KI,  action  slow  and 
incomplete. 

2.  HN03  forms  HI03  and  NO  .     Strong  HN03  must  be  used  (at  least 
sp.  gr.  1.42).     Action  is  slow.     A  very  good  method  of  making  HI03 . 

3.  HH2P02  becomes  H3P04  with  acids  and  with  alkalis. 

4.  H2S  becomes  S  and  HI;  no  action  if  both  substances  be  perfectly  dry 
(Skraup,  (7.  (7.,  1896,  i,  469)  (separation  of  H2S  from  AsH3).     According 
to  Saint-Gilles  (A.  Ch.,  1859,  (3),  57,  221),  in  alkaline  mixture  from  six 
to  seven  per  cent  of  the  sulphur  is  oxidized  to  a  sulphate. 

H2S03  becomes  H2S04  and  HI.  With  a  thiosulphate  a  tetrathionate  is 
formed:  2Na2S203  +  I2  ==  Na2S406  +  2NaI  (Pickering,  J.  (7.,  1880,  37, 
128). 

5.  Cl  becomes  IC1  or  IC13 ,  depending  upon  the  amount  of  chlorine 
present,  water  should  be  absent.     In  the  presence  of  water  HC1  and  HI03 
are  formed;  in  alkaline  mixture  a  chloride  and  a  periodate:  I2  +  ?C1,  + 
16NaOH  =  14NaCl  +  2NaI04  +  8H,0.     HC103  forms  HI03  and  HC1: 
5HC103  +  3I2  +  3H20  ==  6HI03  +  5HC1 . 

6.  Br  becomes  IBr,  decomposed  by  water  (Bornemann,  ^4.,  1877,  189, 
183).     In  alkaline  mixture  with  an  excess  of  Br  a  bromide  and  an  iodate: 
I2  +  5Br2  +  12KOH  =  2KI03  +  lOKBr  +  6H20  .    HBr03   becomes   Br 
andHI03. 

7.  Iodine  combines  with  KI  in  concentrated  solution  to  form  KL(KII2)  . 

7.  Ignition. — See  I. 

8.  Detection. — Iodine  is  recognized  by  the  yellow  to  black  color  when 
mixed  with  water;  the  violet  color  when  dissolved  in  carbon  disulphide; 
the  reddish  color  when  dissolved  in  chloroform  or  ether;  the  blue  color 
when  added  to  a  cold  solution  of  starch;  the  violet  color  of  the  vapors,  etc. 
The  presence  of  tannin  interferes  with  the  usual  tests  for  iodine  unless  a 
drop  or  two  of  ferric  chloride  solution  be  added  (Tessier,  Z.,  1874,  11,  313). 

9.  Estimation. —  (a)  It  is  reduced  to  an  iodide,  precipitated  with  AgN03  ,  and 
after  drying-  at  150°,  weighed  as  Ag"!  .  It  is  estimated  volumetrically  with  a 
standard  solution  of  Na,S,O3  ,  using  starch  as  an  indicator.  (&)  The  iodine 
dissolved  in  potassium  iodide  is  treated  with  an  alkaline  solution  of  hydrogen 
peroxide  in  an  azotometer,  the  oxygen  liberated  being  a  measure  of  the  amount 
of  iodine  present  (Baumann,  Z.  angew.,  1891,  204). 


§280,  5,  HYDRIODIC  ACID.  365 

§280.  Hydriodic  acid.     HI  =127. 928. 
HI-' ,  H  —  I . 

1.  Properties. — Molecular    ircifllit,    127.858.     Vapor    </<m' /'///,    63.927.     A    colorless 
incombustible   gas.     At   atmospheric   pressure    it    solidifies   at  — 51°.     At   0°    it 
liquefies  under  a  pressure  of  3.97  atmospheres  ( Faraday,  A.  Ch.,  1845,   (3),  15, 
266).     The  constant  boiling-  point  of   the   aqueous  solution   of  the  gas  is   127°, 
which  solution  contains  57  per  cent  of  HI  and  has  a  specific  gravity  of  1.694 
(Roscoe,  J.  C.,  1861,  13,  160).     Gaseous  HI  is  dissociated  by  heat,  slowly  at  260°; 
rapidly  at  240°  (Lemoine,  A.  Ch.,  1877,  (5),  12,  145).     Iodine  separates  from  the 
water  solution  of  the  acid  when  exposed  to  the  air. 

2.  Occurrence. — Not  found  free  in  nature,  but  in  combination  as  iodide  or 
iodate. 

3.  Formation. — (a)  By  direct  union  of  the  elements  at  a  full  red  heat  (Merz 
and  Holzmann,  B.,  1889,  22,  869).     (b)  By  direct  union  of  the  elements  in  pres- 
ence of  platinum  black  at  300°  to  400°   (Lemoine,  C.  r.,  1877,  85,  34).     (c)  From 
BaI2  by  adding  HoSO4  in  molecular  proportions,     (d)   By  the  action  of  iodine 
upon  Na2SO3  or  Na2S2O3  (Mene,  C.  r.,  1849,  28,  478).     (e)  By  the  action  of  iodine 
upon  moist  calcium  hypophosphite:  Ca(H2P02)2  +  4I2  +  4H2O  =  CaH4(PO4)2 
+  8HI  (Mene,  I.e.). 

Iodides  are  formed  by  the  direct  action  of  iodine  upon  the  metals;  or  better, 
by'the  action  of  HI  upon  the  oxides,  hydroxides  or  carbonates  of  those  metals 
whose  iodides  are  soluble  in  water.  Iodides  of  lead,  silver  and  mercury  are 
formed  by  precipitation. 

4.  Preparation. — (a)  By  passing  H2S  into  a  mixture  of  finely  divided  iodine 
suspended  in  water,  adding  more  iodine  as  fast  as  the  color  disappears:  2I2  + 
2H2S  =  4HI  +  S2  (Pellagri,  Gazzetta,  1875,  5,  423).     (6)  By  bringing  moist  red 
phosphorus  in  contact  with  iodine:  P4    +   10I2    +   16H2O  =  4H3P'04   +   20HI 
(Meyer,#.,   1887,   20,   3381).     (c)    By   passing   vapors   of   iodine   into   hot   liquid 
paraffine   (Crismer,  B.,  1884,  17,  649).     (d)   By  heating  iodine  with  copaiba  oil 
(Bruylants,  B.,  1879,  12,  2059).     It  cannot  be  prepared  by  adding  H2S04  to  an 
iodide  and  distilling  (5). 

5.  Solubilities. — Iodides  of  lead,  silver,  mercury  and  cuprosum  are  in- 
soluble. Iodides  of  other  ordinary  *  metals  are  soluble,  those  of  bismuth, 
tin  and  antimony  requiring  a  little  free  acid  to  hold  them  in  solution. 
Lead  iodide  is  sparingly  soluble  in  water  (§57,  5c).  Mercuric  iodide  is 
readily  soluble  in  excess  of  potassium  iodide,  forming  a  double  iodide, 
K2HgI4;  most  other  iodides  are  more  soluble  in  a  solution  of  potassium 
iodide  than  in  pure  water.  The  iodides  of  the  alkalis,  Ba ,  Ca  and  Hg" 
are  soluble  in  alcohol;  Hgl  and  Agl  are  insoluble.  All  iodides  in  solution 
are  transposed  by  HC1  or  by  dilute  H2S04 .  Hot  concentrated  H2S04 
decomposes  all  iodides,  those  of  Pb ,  Ag  and  Hg  slowly  but  completely, 
S02  and  I  being  produced:  2KI  +  2H2SO,  =  K2S04  +  I2  +  S02  +  2H20  . 
HN03  in  excess  first  transposes  then  decomposes  soluble  iodides:  6KI  -f- 
8HN03  =  6KN03  +  3I2  +  2NO  +  4H20  .  If  the  HN03  be  concentrated 
the  iodine  is  further  oxidized:  31,  +  10HN03  =  6HI03  +  10NO  +  2H20  . 
Long-continued  boiling  with  HN03 ,  sp.  gr.  1.42,  decomposes  the  insoluble 
iodides.  Chlorine  in  the  cold  decomposes  all  soluble  iodides,  by  heating 
with  chlorine  the  insoluble  iodides  are  also  decomposed:  2KI  +  C12  = 

*  Thallium  iodide,  Tl  I,  is  perfectly  insoluble  in  cold  water,  a  distinction  and  separation  froiij 
bromides  and  chlorides  (Huebner,  Zf,  1872,  II,  397).  Palladous  iodide  is  insoluble  in  water. 


366  HTDRIODIC  ACID.  §280,  6A. 

2KC1  -f-  ^2  •  With  an  excess  of  chlorine  the  iodine  is  further  oxidized : 
I2  +  5C12  +  6H20  =  2HI03  -f  10HC1 .  Silver  iodide  is  almost  insoluble 
in  ammonium  hydroxide  or  ammonium  carbonate  (distinction  from  silver 
chloride).  It  is  soluble  in  KCN .  Agl  and  PbI2  are  soluble  by  decomposi- 
tion in  solution  of  alkali  thiosulphates :  Agl  -f-  Na2S203  =  Nal  -f- 
NaAgS203 .  Lead  iodide  is  soluble  in  a  solution  of  the  fixed  alkalis. 

6.  Reactions. — A. — With  metals  and  their  compounds. — Silver  nitrate 
solution  in  excess  precipitates,  from  solutions  of  iodides,  silver  iodide,  Agl , 
yellow-white,  blackening  in  the  light  without  appreciable  separation  of 
iodine.  For  solubilities  see  paragraph  above. 

Solution  of  mercuric  chloride  precipitates  the  bright,  yellowish-red  to 
red,  mercuric  iodide,  HgI2 .  The  precipitate  redissolves  on  stirring,  after 
slight  additions  of  the  mercuric  salt,  until  equivalent  proportions  are 
reached,  when  its  color  deepens.  For  the  solubilities  of  the  precipitate 
see  §58,  6/.  Solution  of  mercurous  nitrate  precipitates  mercurous  iodide, 
Hgl ,  yellow  to  green  (§58,  6/). 

Solution  of  lead  nitrate  or  acetate  precipitates,  from  solutions  of  iodides 
not  very  dilute,  lead  iodide,  PbI2 ,  bright-yellow — soluble,  as  stated  in  full 
in  §57,  5c. 

Palladous  chloride,  PdCl2,  precipitates,  from  solutions  of  iodides,  pal- 
ladous  iodide,  PdI2 ,  black,  insoluble  in  water,  alcohol  or  dilute  acids,  and 
visible  in  500,000  parts  of  solution.  The  reagent  does  not  precipitate 
bromine  at  all  in  moderately  dilute  solutions,  slightly  acidulated  with  HC1 . 
Palladous  iodide  is  slightly  soluble  in  excess  of  the  alkali  iodides,  and  is 
soluble  in  ammonium  hydroxide  (§106). 

Copper  salts  precipitate  from  solutions  of  iodides  cuprous  iodide  (white) 
mixed  with  iodine  (black) :  2CuS04  +  4KI  ==  2CuI  +  2K,S04  +  I2 .  If 
sufficient  reducing  agents  (as  sulphurous  acid)  are  present  to  reduce  the 
liberated  iodine  to  HI ,  only  the  white  cuprous  iodide  will  be  precipitated 
(a  distinction  from  bromides  and  chlorides). 

When  metals  are  attacked  by  HI  an  iodide  is  formed  and  hydrogen  is 
evolved.  Hydriodic  acid  unites  with  all  metallic  oxides  and  hydroxides 
(expect  ignited  Cr203)  to  form  iodides;  frequently,  however,  iodine  is 
liberated  and  an  iodide  of  lower  metallic  valence  is  formed: 

1.  Pb"+n  becomes  Pb"  . 

2.  Asv  becomes  As"' ;  KI  has  no  action  upon  normal  K3As04  (Friedheim 
and  Meyer,  Z.  anorg.,  1891,  1,  409). 

3.  Sbv  becomes  SV"  . 

4.  Biv  becomes  Bi'"  . 

5.  Cu"  becomes  Cu' .     Soluble  iodides  reduce  normal  cupric  salts,  but 
have  no  reducing  action  in  alkaline  mixture  or  upon  cupric  hydroxide. 
With  phenylhydrazine  sulphate  and  cupric  sulphate  the  iodine  of  iodides  is 


§280,  6£,  6.  HYDRIODIC  ACID.  367 

completely  precipitated  (separation  from  chlorides)  (Kaikow,  Ch.  Z.,  1894, 
18,  1661). 

6.  Fe/r/  becomes  Fe"  (§269,  8). 

7.  CrVI  becomes  Cr'" .     K2Cr04  is  not  reduced  by  KI  even  upon  boiling 
the  concentrated  solutions.     K2Cr207  with  KI  slowly  gives  I  and  Cr'"  in 
the  cold.     When  KI  is  boiled  with  a  concentrated  solution  of  K2Cr207  the 
iodine  is  completely  liberated  (separation  from  bromides  and  chlorides 
which  are  unchanged):   6KI  +   5K2Cr207  =  8K,Cr04   +   Cr203   +   3I2 
(Dechan,  J.  C.,  1886,  50,  682;  1887/51,  690).     When  Agl  is  boiled  with 
K2Cr207  and  H2S04  no  iodine  is  evolved,  chromium  is  reduced  and  the 
iodide  becomes  silver  iodate:  K2Cr207  +  Agl  +  5H2S04  =  2KHS04  + 
Cr2(S04)3  +  AgI03  +  4H20  (Macnair,  J.  C.,  1893,  63,  1051). 

8.  Co"+n  becomes  Co";  KI  has  no  reducing  action  upon  cobaltic  hy- 
droxide. 

9.  Ni"+n  becomes  Ni";  KI  reduces  Ni'" ,  liberating  iodine. 

10.  Mn"+n  becomes  Mn"  .     When  KI  is  boiled  with  KMn04  the  manga- 
nese becomes  Mn02    and  the  iodide  is  oxidized  to  an  iodate:  6KMn04  + 
3KI  +  .3H20  =  3KI03  +  6Mn02  +  6KOH  (Groeger,  Z.  angew.,  1894,  13 
and  52)  (distinction  from  bromides,  which  do  not  decolor  permanganates). 

B. — With  non-metals  and  their  compounds. 

1.  H3Fe(CN)G  forms  H4Fe(CN)6  and  I;  the  reaction  also  takes  place  in 
neutral  mixture. 

2.  HN02  forms   NO  and  I   (separation   of   iodide   from  bromide   and 
chloride)  (Jannasch  and  Aschoff,  Z.  anorg.,  1891,  1,  144  and  245). 

HN03  forms  NO  and  I,  with  further  oxidations  to  HIO;5  with  concen- 
trated HN03 .     The  HN02  acts  much  more  rapidly  than  the  HN03 . 

3.  No  reduction  with  phosphorous  compounds. 

4.  H2S04  dilute  no  action;  with  the  concentrated  acid  in  excess,  S02  and 
I  are  formed:   2KI  +  3H2S04  =  I2  +  S02  +  2KHS04  +  2H20  ;  if  KI  be 
added  in  excess  to  boiling  H2S04  ,  H2S  and  I  are  formed:   SKI  +  9H2S04  = 
4I2  +  H2S  +  8KHS04  +  4H20  (Jackson,  J.  (7.,  1883,  43,  339).     Ammo- 
nium persulphate  liberates  iodine  from  iodides  at  ordinary  temperature 
(Engel,  C.  r.,  1894,  118,  1263). 

5.  Cl  in  excess  forms  HC1  and  HI03 ;  with  excess  of  HI ,  HC1  and  I  are 
formed.     In  the  presence  of  a  fixed  alkali  a  periodate  and  a  chloride  are 
formed:   KI  +  8KOH  +  4C12  =  8KC1  +  KI04  +  4H20 .     Hypochlorous 
acid  oxidizes  to  iodine,  then  to  iodic  in  acid  solution;  in  alkaline  solution 
to  periodate. 

HC103  with  excess  of  HI  forms  HC1  and  I;  with  excess  of  HC103  HC1 
and  HI03 . 

6.  Br  forms  I  and  HBr  or  a  bromide. 


368  BYDRIODIC  ACID.  §280, 6B,  7. 

HBr03  with  excess  of  HI  forms  HBr  and  I ;  with,  excess  of  HBr03 ,  Br 
and  HI03 . 

7.  HI03,  iodine  is  liberated  from  both  acids:  HI03  +  5HI  =  3I2  -(- 
3H20  .     HI04  gives  iodine. 

8.  H202  becomes  H20 ,  0  and  I  (§244,  QB6)  (Cook,  /.  C.,  1885,  47,  471). 

9.  Ozone  promptly  liberates  iodine  from  soluble  iodides.     Atmospheric 
oxygen  decomposes  HI  and  ferrous  and  calcium  iodides  slowly,  the  alkali 
iodides  not  at  all. 

7.  Ignition — As  a  general  rule  iodides  strongly  ignited  in  presence  of  air 
and  moisture  evolve  iodine,  leaving  the  oxide  of  the  metal.  Ignited  in  absence 
of  air  or  moisture  the  following  iodides  are  not  decomposed:  KI ,  Nal  ,  BaI2  , 
CaI2  ,  SrI2  ,  Mnl,  ,  A1I3  ,  SnI4  ,  PbI2  ,  Agl  and  Hgl,  .  See  Mitscherlich  (Pogg., 
1833,  29,  193),  Personne  (C.  r.,  1862,  54,  216)  and  Gustavson  (A.,  1873,  172,  173). 

8.  Detection. — The  iodide  is  oxidized  to  free  iodine  by  one  of  the  re- 
agents mentioned  in  (6)  above.  With  a  dry  powder  hot  concentrated 
H2S04  is  usually  employed  when  the  iodine  is  detected  by  the  violet  fumes 
evolved,  condensing  in  the  cooler  portion  of  the  test  tube.  With  solu- 
tions the  usual  reagent  is  chlorine  water.  The  iodine  is  recognized  by 
the  violet  color  when  shaken  with  CS2 ,  or  the  bright-red  color  with  CHC13 . 
In  case  a  large  amount  of  iodine  be  present  the  CS2  solution  may  be  almost 
black.  In  this  case  large  dilution  with  CS2  is  necessary  to  detect  the  violet 
color.  If  but  a  small  amount  of  iodine  be  present  the  chlorine  must  be 
added  very  cautiously  or  the  iodide  will  all  be  oxidized  to  the  colorless 
iodic  acid.*  With  small  amounts  of  iodide,  nitric  acid  is  less  liable  to 
cause  error  as  relatively  much  more  nitric  acid  is  required  to  oxidize  the 
iodine  to  iodic  acid.  For  the  detection  of  small  amounts  of  iodide  a 
cupric  salt  strongly  acidulated  with  HC1  is  an  excellent  reagent  for  the 
oxidation :  2CuCl2  +  2KI  =  2CuCl  +  2KC1  +  I2 .  The  addition  of  HN02 
or  NaN02  to  the  acid  solution  will  also  liberate  iodine  but  not  bromine 
or  chlorine.  The  iodine  may  then  be  detected  and  removed  by  shaking 
with  CS2 .  Bromine  and  chlorine  may  then  be  tested  for. 

If  insoluble  iodides  are  present  they  should  be  transposed  by  H2S , 
the  insoluble  sulphide  removed  by  nitration,  the  excess  of  H2S  removed 
by  boiling,  and  the  solution  then  tested  for  hydriodic  acid.  Or  the 
insoluble  iodide  should  be  reduced  by  Zn  and  H2S04 :  2AgI  +  Zn  +  H2S04 
=  2Ag  -f  ZnSO,  +  2HI .  The  nitrate  may  then  be  tested  for  hydriodic 
acid.  The  insoluble  iodide  may  also  be  fused  with  NaoG03 ,  and  after 
digestion  with  water  the  filtrate  acidulated  and  tested  for  hydriodic  acid. 
That  is,  the  solution  must  be  acidulated  before  chlorine  water  is  added, 
else  the  iodine  will  be  oxidized  to  an  iodate  or  periodate. 

*The  test  potassium  bromide  for  traces  of  an  iodide  it  is  recommended  to  add  CSj  and  cupric 
sulphate  or  a  small  amount  of  ferric  alum.  Or  add  chlorine  water  and  then  a  few  crystals  of 
ferrous  sulphate;  then  shake  with  CS2  (Brito,  C.  N.,  1884,  50,  210). 


§281,  5.  WDIC  ACID.  369 

9.  Estimation. — Gravimetrically  by  precipitation  as  Agl  and  weighing  as  such 
after  gentle  ignition.  Volumetrically  by  oxidation  to  iodine  and  titration  with 
standard  Na2S2O3  (Groger,  Z.  angew.,  1894,  52). 

§281.  lodic  acid.     HI03  =  175.928. 
H'IvO-"3,  H  —  0  —  I  ~ 

1.  Properties. — lodic  acid  is  a  white  crystalline  solid;  its  solution  saturated 
at  14°  contains  (is. 5  per  cent  HIO3  ,  and  has  a  specific  gravity  of  2.1629  (Kaem- 
merer,  Poyy.,  18G(J,  138,  390).     At  170°  it  loses  water,  forming  iodic  anhydride, 
I2O5  ,    u   white    crystalline   solid,    which,    at    300°,    dissociates   into    iodine    and 
oxygen.     See  Ditte,  A.  Ch., .1870,   (4),  21,  5.     It  is  readily  soluble  in  water  and 
in  alcohol;  the  solutions  redden  litmus  and  afterwards  bleach  it. 

2.  Occurrence. — The  free  acid  is  not  found  in  nature.     It  is  found  as  Ca(IO3)2 
in  sea  water,  and  us  sodium  iodate  in  Chili  saltpeter  (Sonstadt,  C.  N.,  1872,  25, 
196,  231  and  241;  Guyard,  BL,  1874,  (2),  22,  60). 

3.  Formation. —  (a)  By  electrolyzing  a  solution  of  I  or  HI  (Riche,  C.  r.,  18">8, 
46,   348).     (b)    By  the   action  of  chlorine   on  iodine  in   the  presence   of   much 
water.     The  HC1  formed  cannot  be  expelled  by  boiling  without  decomposing 
the   HIO3  .     It   must   be   removed   by   the   careful   addition    of   Ag20  .     (c)    By 
adding  water  to  IC13    and  wrashing  with  alcohol:   2IC13    -4-    3H2O  =  HIO3    + 
5HC1  +  IC1 .     ((/)  KIO3  is  made  by  treating  iodine  with  KOH:  3la  +  6KOH  = 
KIOj  +  r>KI  +  3H2O  .     And  then  washing  with  alcohol  to  remove  the  KE  .     (e) 
By  heating  potassium  chlorate  and  iodine:  10KC103  +  6I2  +  6H2O  :==  6KHI2O6 
+  4KC1  +  6HC1  (Bassett,  J.  C.,  1890,  57,  760).   (f)  By  boiling  iodine  with  barium 
hydroxide  until  neutral,  filtering  and  decomposing  with  sulphuric  acid  (Steven- 
son, C.  N.,  1877,  36,  201).     (g)  By  the  action  of  I  upon  AgNO3:  5Ag-N03  +  31,  + 
3H,0  =  5 Agl  +  5HN03  +  HIOa  . 

lodates  of  the  alkalis  and  alkaline  earths  are  easily  made  by  the  action  of 
iodine  on  the  hydroxides,  and  separation  by  alcohol  or  by  crystallization  from 
the  iodides  which  are  formed  in  the  reaction.  All  iodates  may  be  made  by 
action  of  the  acid  on  the  hydroxides  or  carbonates. 

4.  Preparation. —  (a)    Iodine   is   oxidized   by   boiling  with   nitric   acid  sp.   gr. 
1.52,  and  removing  the  excess  of  the  nitric  acid  by  evaporation.     (6)  By  adding 
a  slight  excess  of  H2S04   to  Ba(IO3)2  and  removal  of  the  excess  of  H2SO4  by 
the    careful   addition    of    Ba(IOi)2  •     (<?)    By    boiling    a    solution    of    potassium 
iodide  wilh  an  excess  of  potassium  permanganate  in  neutral  or  alkaline  solu- 
tion: KI  +  2KMn04   +  H,O  —  KIO3   +  2KOH  +  2MnO2    (Groger,  Z.  angew., 
1894,  13  and  52).     (<1)  The  very  stable  potassium  biiodate,  KHI2O8  ,  is  formed  by 
recrystallizing  a  water  solution   of  equal  portions  of  KIO3    and  HI03  .     It  is 
soluble  in  18.66  parts  water  at  17°  (Meineke,  A.,  1891,  261,  359). 

5.  Solubilities. — Ba(I03)2  is  soluble  in  about  3000  parts  water  at  ordi- 
nary temperature;  and  in  about  GOO  parts  at  100°  (Kremers,  Pogg.,  1851, 
84,  27;  Spica,-  Gazzetia,  1894,  24,  i,  91).  AgI0.5  is  soluble  in  27/700  parts 
of  water  at  25°;  in  2.1  parts  NH4OH  (10  per  cent)  at  25°  (separation  from 
silver  iodide);  in  1044.3  parts  HN03 ,  sp.  gr.  1.21  at  25°  (Longi,  Gazzetta, 
1883,  13,  87).  The  iodatos  of  Ag  ,  Ba  ,  Pb  ,  Hg  ,  Sn  ,  Bi  ,  Cd  ,  Fe  and  Cr 
require  at  15°  more  than  500  parts  of  water  for  their  solution  and  the 
following  require  less :  Cu  ,  Al ,  Co  ,  Ni ,  Mn  ,  Zn  ,  Sr ,  Ca ,  Mg ,  K  and  Na  . 
They  are  all  transposed  by  concentrated  HN03  or  H2S04;  and  are  decom- 
posed by  concentrated  HC1 .  They  are  soluble  in  the  alkalis  in  so  far  as 


370  IODIC  ACID.  §281,  6. 

the  corresponding  metallic  oxides  are  soluble  in  those  reagents.  Most 
of  the  iodates  are  insoluble  in  alcohol  (with  K,  Na,  Ba  and  Ca  iodates  a 
separation  from  iodides). 

6.  Reactions. — A. — With  the  metals  and  their  compounds. — A  few  metals 
are  attacked  evolving  hydrogen,  forming  iodates,  sometimes  traces  of 
iodides.  With  the  following  metallic  compounds  the  valence  of  the  metal 
is  changed: 

1.  As'"  becomes  Asv  with  liberation  of  iodine.     AsH3  in  excess  forms 
As0  ,  with  the  HI03  in  excess  Asv  (Ditte,  A.,  1870,  156,  336). 

2.  Sb'"  becomes  Sbv  with  liberation  of  iodine.     SbH3  forms  Sb°  . 

3.  Sn"  becomes  SnIV  and  HI . 

4.  Cu'  becomes  Cu"  with  liberation  of  iodine. 

5.  Fe"  becomes  Fe'"  with  liberation  of  iodine. 

Solution  of  silver  nitrate  precipitates,  from  even  very  dilute  solutions  of 
iodates  and  from  solutions  of  iodic  acid  if  not  very  dilute,  silver  iodate, 
AgI03 ,  white,  crystalline,  soluble  in  ammonium  hydroxide,  soluble  in  an 
excess  of  hot  HN03 .  In  the  ammoniacal  solution,  hydrosulphuric  acid 
forms  silver  sulphide,  sulphur  and  ammonium  iodide. 

Barium  chloride  precipitates  barium  iodate,  Ba(I03)2 ,  slightly  soluble 
in  cold,  more  soluble  in  hot  water,  insoluble  in  alcohol,  soluble  in 
hot  dilute  nitric  acid,  readily  soluble  in  cold  dilute  hydrochloric  acid. 
Hence,  dilute  solutions  of  free  iodic  acid  should  either  be  neutralized  or 
tested  with  barium  nitrate.  This  precipitate,  by  addition  of  alcohol,  is  a 
complete  separation  from  iodides,  and,  when  well  washed,  decomposed  with 
a  very  little  sulphurous  acid  (8),  and  found  to  color  carbon  disulphide 
violet,  its  evidence  for  iodic  acid  is  conclusive.  Barium  iodate  is  trans- 
posed by  ammonium  carbonate. 

Salts  of  lead  give  a  white  precipitate  of  lead  iodate,  Pb(I03)2 .  Ferric 
chloride  gives,  in  solutions  not  dilute,  a  yellowish-white  precipitate  of 
ferric  iodate,  Fe(I03)3,  sparingly  soluble  in  water,  and  freely  soluble  in 
excess  of  the  reagent.  Boiling  decomposes  it. 

Alcohol  precipitates  potassium  iodate  from  water  solution,  an  approxi- 
mate separation  from  iodide. 

B. — With  non-metals  and  their  compounds. 

1.  H2C204  becomes  C02  and  I .     Action  is  slow  unless  solutions  are  hot. 
Carbon  (except  diamond)  heated  in  sealed  tubes  becomes  C02  with  sepa- 
ration of  I  (Ditte,  I.  c.). 

H4Fe(CN)6  becomes  H,Fe(CN)6  and  I . 

HCNS  forms  H2S04 ,  I  and  some  other  products. 

2.  HN02  becomes  HN03  and  I . 

3.  PH3  becomes  H3P04  and  I .     With  an  excess  of  PH3 ,  HI  is  formed. 


§281,  9c.  IODIC  ACID.  371 

Water  in  which  phosphorus  has  stood  reduces  iodic-acid  to  iodine  (Corne, 
J.  Pharm.,  1878,  (4),  28,  386). 
HHJ?0,  becomes  H3P04  and  I. 

4.  H,S  becomes  S  and  I.    Thiosulphates  form  first  iodine,  then  an  iodide. 
H2S03 ,  with  excess  of  HI03 ,  becomes  H2S04  and  I;  with  excess  of  H2S03 , 

H2S04  and  HI . 

5.  HC1 ,  if  concentrated,  forms  IC13  and  Cl ,  iodine  not  being  liberated. 
(>.  HBr  forms  Br  and  I . 

7.  HI  forms  I  from  both  acids.     The  addition  of  tartaric  acid  to  a  mix- 
ture of  KI  and  KI03  is  sufficient  to  give  an  immediate  test  for  free  iodine 
with  CS2 .     It  must  be  remembered  that  an  iodide  alone  rendered  acid  will 
give  a  test  for  free  iodine  after  a  short  time. 

8.  Morphine  reduces  iodic  acid  with  separation  of  iodine. 

7.  Ignition. — Potassium  and  sodium  iodates  on  ignition  form  iodides 
and  evolve  oxygen  (Cook,  J.  C.,  1894,  65,  802).     Many  other  iodates  evolve 
oxygen  but  the  iodide  formed  is  further  decomposed  as  stated  in  §275,  7. 

Iodates  in  dry  mixture  with  combustible  bodies  are  reduced,  on  heating 
or  concussion,  with  detonation,  but  much  less  violently  than  chlorates  or 
nitrates. 

8.  Detection. — It  is  usually  detected,  after  acidulation,  by  treatment 
with  some  reducing  agent  for  the  formation  of  free  iodine.     H2S03  is 
often  employed  because  it  acts  rapidly  and  in  the  cold;  but  traces  of  HI03 
frequently  escape  detection  for  the  least  excess  of  H2S03  at  once  reduces 
the  iodine  to  colorless  hydriodic  acid.     A  desirable  reagent  for  this  reduc- 
tion is  one  that  will  act  rapidly  in  the  cold,  and  in  no  case  cause  the 
further  reduction  to  hydriodic  acid.     The  following  reducing  agents  have 
been  used :  K4Fe(CN)B  acidulated  with  dilute  H2S04 ,  H3As03 ,  Cud  ,  FeS04 , 
morphine  sulphate  and  uric  acid.    -To  detect  KI03  in  KI  it  is  recom- 
mended by  Schering  (J.  (7.,  1873,  26,  191)  to  add  a  crystal  of  tartaric 
acid  to  the  solution.     The  formation  of  a  yellow  zone  is  indicative  of  an 
iodate.     Hydrochloric   acid  may  be  used,  but  if  it  contains  a  trace  of 
chlorine  it  will  give  the  test  for  an  iodate.     Iodine  frequently  occurs  in 
nitric  acid  as  iodic  acid.     Hilzer  (J.  (7.,  1876,  29,  442)  directs  io  add  equal 
volumes  of  water,  carbon  disulphide,  and  then  coarse  zinc  filings.     It  may 
be  necessary  to  warm  the  solution  slightly.     Biltz  (C.  C.,  1877,  Sfi)  dilutes 
the  HNOr>  with  water,  adds  starch  solution  and  then  H2S  solution  drop 
by  drop.     A  blue  zone  is  formed  if  HI03  be  present. 

0.  Estimation.— (a)  B.v  precipitation  with  AgN03  ,  and  after  drying  at  100° 
weighing-  as  Ag-10,.  (ft)  By  reducing-  to  an  iodide  and  estimating-  as  snch. 
(r)  By  treating-  with  KI  acidulated  with  H2SO,  ,  and  titrating  the  iodine  lib- 
erated with  standard  Na2S2O,  . 


372  PERIODIC  ACID.  §282. 

§282.  Periodic  acid.     HI04  =  191.928. 

H       H      H 
\      I      / 
0  000 

II  XI  / 

H'IVII0-"4  or  H'5IVII0-"6 ,  H  —  0  — 1  =  0  orH  —  0  —  I  —  0  —  H. 


The  anhydride,  I207  ,  has  not  been  isolated,  and  but  one  acid  is  known  in  the 
free  condition,  HIO4,2H2O  or  H5IO6  .  This  acid  exists  in  colorless  monoclinic 
crystals,  which  do  not  lose  water  at  100°.  It  melts  at  133°,  and  at  a 
higher  temperature  it  decomposes  into  iodic  anhydride,  water  and  oxygen 
(Kimmins,  J.  C.,  1887,  51,  356;  and  1889,  55,  148).  Numerous  periodates  have 
been  prepared  as  if  derived  from  one  or  the  other  following1  named  acids: 
HI04  ,  H3I05  ,  H3I06  ,  H4I209  ,  HJ.O^  ,  H12I2013  ,  H10I4O19  ,  H10I6026 
(Rammelsberg,  Pogg.,  1865,  134,  368,  499). 

The  free  periodic  acid,  H5IOC  ,  is  prepared:  (a)  By  oxidizing  iodine  with  per- 
chloric acid:  2HC1O4  +  I2  +  4H,0  =  2Hr>I06  +  C12  (Kaemmerer,  Pogg.,  1869, 
138,  406).  (I))  By  heating  iodine  or  barium  iodide  with  a  mixture  of  barium 
oxide  and  barium  peroxide,  digesting  with  water,  and  transposing  the 
Ba5(I06)2  thus  obtained  with  the  calculated  amount  of  sulphuric  acid  (Ram- 
melsberg, Pogg.,  1869,  137,  305).  (c)  By  conducting  chlorine  into  sodium  iodate 
in  presence  of  sodium  hydroxide:  NalO,  +  3NaOH  -f  C12  =  Na2H8IOe  + 
2NaCl  .  This  acid  periodate  dissolved  in  water  with  a  little  nitric  acid  and 
then  precipitated  with  silver  nitrate,  forms  the  silver  salt,  AgoH-IOe  .  This 
precipitate  is  dissolved  in  nitric  acid  and  evaporated  on  the  water-bath,  when 
orange-colored  crystals  of  silver  met  a  periodate  are  formed  according  to  the 
following:  2Ag2H3'lO6  +  2HN03  =  2AgI04  +  2AgN03  +  4H2O  .  Water  decom- 
poses this  precipitate:  2AgIO4  +  4H2O  =  H3IO6  +  Ag2H3I06  .  Or  the  silver 
periodate,  AgI04  ,  is  decomposed  by  Cl  or  Br  (Kaemmerer,  ?.  c.,  p.  390). 

The  silver  salts  vary  in  color:  AgI04  is  orange:  Ag2HIO5  ,  dark  brown; 
Ag4I209  ,  chocolate  colored;  while  silver  iodate  is  white  (a  distinction).  In  the 
general  reactions  periodic  acid  and  periodates  resemble  iodic  acid  and  iodates. 

H2C204  becomes  CO2  and  I . 

H3PO2  becomes  H3P04  and  HI . 

H2S  becomes  S  and  HI . 

H2S03  becomes  H2SO4  and  HI03  without  separation  of  iodine  when  the  two 
acids  are  present  in  molecular  proportions.  The  presence  of  a  greater  pro- 
portion of  H,SO3  causes,  first,  separation  of  iodine  with  final  complete  reduc- 
tion to  HI  (Selmous,  B.,  1888,  21,  230): 

HI04  +    H2S03  =  HI03  +  H2S04 
3HI04  +  8H2S03  =  HI03  -f  I2  +  8H,SO«  +  H2O 
2HI04  +  7H2S03  =  I3  +  7H,S04  +  H2O 
HI04  +  4H2S03  =  HI  +  4H2S04 

HC1  becomes  Cl  and  IC18 

HI  forms  I  from  both  acids. 

According  to  Lautsch  (J.  pr.,  1867,  100,  86),  its  behavior  with  mercurons 
nitrate  is  characteristic.  The  pentasodic  periodate,  Na5IOa  ,  gives  a  light- 
yellow  precipitate,  Hg5IO,  , 


§283.         COMPARATIVE  REACTIONS  OF  HALOGEN  COMPOUNDS.  373 


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PART  IV -SYSTEMATIC  EXAMINATIONS. 


REMOVAL    OF   ORGANIC  SUBSTANCES. 

L  Before  the  liquid  reagents  can  be  applied,  solids  must  be  dissolved. 
The  methods  of  inorganic  analysis  do  not  provide  against  interference  by  organic 
compounds;  and,  in  general,  it  is  impossible  to  conduct  inorganic  analysis  in 
material  containing  organic  bodies.  It  is,  therefore,  first  necessary  to  ascertain 
if  the  unknown  substance  contains  organic  matter.  This  has  been  ascertained  in 
the  preliminary  examination  by  the  blackening  or  charging  of  the  substance  when 
heated  in  the  closed  tube.  A  burnt  odor  is  also  excellent  evidence  of  the  presence 
of  organic  matter.  If  organic  matter  is  found  to  be  present,  it  may  be  removed  as 
follows:  1st,  by  combustion  at  a  red  or  white  heat,  with  or  without  oxidizing 
reagents;  2d  (in  part),  by  oxidation  with  potassium  chlorate  and  hydrochloric 
acid  on  the  water-bath  (§69,  6' 'el);  3d,  by  oxidation  with  nitric  acid  in  presence 
of  sulphuric  acid,  at  a  final  temperature  of  the  boiling  point  of  the  latter  (§79,  6'e3); 
4th,  by  solvents  of  certain  classes  of  organic  substances;  5th,  by  dialysis.  These 
operations  are  conducted  as  follows: 

§285.  Combustion  at  a  red  or  white  heat,  of  course,  excludes  analysis  for  mer- 
cury, arsenous  and  antimonous  bodies  (except  as  provided  in  '§70,  7),  and 
ammonium.  The  last-named  constituent  can  be  identified  from  a  portion  of  the 
material  in  presence  of  the  organic  matter  (§207,  3).  If  chlorides  are  present 
some  iron  will  be  lost  at  temperatures  above  100°,  and  potassium  and  sodium 
waste  notably  at  a  white  heat,  and  slightly  at  a  full  red  heat.  Certain  acids 
will  be  expelled,  and  oxidizing  agents  reduced. 

The  material  is  thoroughly  dried  and  then  heated  in  a  porcelain  or  platinum 
crucible,  at  first  gently.  It  will  blacken,  by  separation  of  the  carbon  of  the 
organic  compounds.  The  ignition  is  continued  until  the  black  color  of  the 
carbon  has  disappeared.  In  special  cases  of  analysis,  it  is  only  necessary  to 
char  the  material;  then  pulverize  it,  digest  with  the  suitable  solvents,  and 
filter;  but  this  method  does  not  give  assurance  of  full  separation  of  all  sub- 
stances. Complete  combustion,  without  use  of  oxidizing  agents,  is  the  way 
most  secure  against  loss,  and  entailing  least  change  of  the  material;  it  is,  how- 
ever, sometimes  very  slow.  The  operation  may  be  hastened,  with  oxidation  of 
all  materials,  by  addition  of  nitric  acid,  or  of  ammonium  nitrate.  The  material  is 
first  fully  charred;  then  allowed  to  cool  till  the  finger  can  be  held  on  the 
crucible;  enough  nitric  acid  to  moisten  the  mass  is  dropped  from  a  glass  rod 
upon  it,  and  the  heat  of  the  water-bath  continued  until  the  mass  is  dry,  when 
it  may  be  very  gradually  raised  to  full  heat.  This  addition  may  be  repeated 
as  necessary.  The  ammonium  nitrate  may  be  added,  as  a  solid,  in  the  same 
way.  The  residue  is  treated  according  to  §301. 

§286.  Oxidation  with  potassium  chlorate  and  hydrochloric  acid  on  the  icater-bath 
does  not  wholly  remove  organic  matter,  but  so  far  disintegrates  and  changes 
it  that  the  filtrate  will  give  the  group  precipitates,  pure  enough  for  most  tests. 
It  does  not  vaporize  any  bases  but  ammonium,  but  of  course  oxidizes  or 
chlorinates  all  constituents.  It  is  especially  applicable  to  viscid  liquids;  it  may 
be  followed  by  evaporation  to  dryness  and  ignition,  according  to  the  paragraph 
above. 

The  material  with  about  an  equal  portion  of  hydrochloric  acid  is  warmed  on 
the  water-bath,  and  a  minute  portion  of  potassium  chlorate  is  added  at  short 
intervals,  stirring  with  a  glass  rod.  This  is  continued  until  the  mixture  is 
wholly  decolored  and  dissolved.  It  is  then  evaporated  to  remove  chlorine, 
diluted  and  filtered.  If  potassium  and  chlorine  are  to  be  tested  for,  another 
portion  may  be  treated  with  nitric  acid,  on  the  water-bath.  The  organic 
matter  left  from  the  action  of  the  chlorine  or  the  nitric  acid  may  be  sufficient 
to  prevent  the  precipitation  of  aluminum  and  chromium  in  the  third  group  of 
bases;  so  that  a  portion  must  be  ignited.  As  to  arsenic  and  antimony,  see 
§70,  7. 

§287.  The  action  of  sulphuric  irith  nitric  acid  at  a  gradually  increasing  heat 
leaves  behind  all  the  metals  (not  ammonium),  with  some  loss  of  mercury  and 
arsenic  (and  iron?)  if  chlorides  are  present  in  considerable  quantity.  In  this, 
as  in  the  operations  before  mentioned,  volatile  acids  are  lost — sulphides  partly 
oxidized  to  sulphates,  etc. 

The  substance  is  placed  in  a  tubulated  retort,  with  about  four  parts  of  con- 
centrated sulphuric  acid,  and  gently  heated  until  dissolved  or  mixed.  A  funnel 


§292.  PRELIMINARY   EXAMINATION  OF  SOLIDS.  375 

is  now  placed  in  the  tubulure,  and  nitric  acid  added  in  small  portions,  gntdu- 
ally  raising  the  heat,  for  about  half  an  hour — so  as  to  expel  the  chlorine,  and 
n  t  vaporize  chlorides.  The  material  is  now  transferred  to  a  platinum  dish 
and  heated  until  the  sulphuric  acid  begins  to  vaporize.  Then  add  small  portions 
of  nitric  acid,  at  intervals,  until  the  liquid  ceases  to  darken  by  digestion,  after 
a  portion  of  nitric  acid  is  expelled.  Finally,  evaporate  off  nearly  all  of  the  sulphuric 
acid,  using  the  lowest  possible  heat  at  the  close.  Cool  and  dilute  with  a  few  c.c. 
of  water.  If  a  residue  remains  undissolved  boil  for  a  few  minutes,  cool,  filter  and 
wash  with  dilute  sulphuric  acid.  The  residue  will  contain  (a)  undecomposed 
mineral  matter  such  as  silicates,  (6)  all  the  lead,  strontium  and  barium  as  sul- 
phates, (c)  some  of  the  calcium,  bismuth,  antimony  and  tin  and  (d)  practically  all  of 
the  chromium  as  the  insoluble  anhydrous  sulphate. 

§288.  The  solvents  used  are  chiefly  ether  for  fatty  matter,  and  alcohol  or  ether, 
or  both  successively,  for  resins.  Instead  of  either  of  these,  benzol  may  be 
used:  and  many  fats  and  some  resins  may  be  dissolved  in  petroleum  ether. 
It  will  be  observed  that  ether  dissolves  some  metallic  chlorides,  and  that 
alcohol  dissolves  various  metallic  salts.  Before  the  use  of  either  of  these  sol- 
vents upon  solid  material,  it  should  be  thoroughly  dried  and  pulverized.  Fatty 
matter  suspended  in  water  solutions  may  be  approximately  removed  by  filter- 
ing- through  wet,  close  filters;  also  by  shaking-  with  ether  or  benzol,  and  decant- 
ing- the  solvent  after  its  separation. 

§289.  H\i  mall/six,  the  larger  part  of  any  ordinary  inorganic  substance  can 
be  extracted  in  approximate  purity  from  the  greater  number  of  organic  sub- 
stances in  water  solution.  The  degree  of  purity  of  the  separated  substance 
depends  upon  the  kind  of  organic  material.  Thins  albuminoid  compounds  are 
almost  fully  rejected;  but  saccharine  compounds  pass  through  the  membrane 
quite  as  freely  as  some  metallic  salts.  (Consult  Watts'  Dictionary,  1894,  IV,  172). 

PKELIMINAEY  EXAMINATION  OF  SOLIDS.* 

§290.  Before  proceeding  to  the  analysis  of  a  substance  in  the  wet  way,  a 
careful  study  should  usually  be  made  of  the  reactions  which  the  substance 
undergoes  in  the  solid  state,  when  subjected  to  a  high  heat,  either  alone  or  in 
the  presence  of  certain  reagents,  before  the  blow-pipe,  or  in  the  flame  of  the 
Bunsen  burner.  This  examination  in  the  dry  way  precedes  that  in  the  wet, 
and  should  be  carried  on  systematically,  following  the  plan  laid  down  in  the 
tables,  and  noting  carefully  every  change  which  the  substance  under  investiga- 
tion undergoes,  and  if  necessary  making  reference  to  some  of  the  standard 
works  on  blow-pipe  analysis.  In  order  to  understand  fully  the  nature  of  these 
reactions,  the  student  should  first  acquaint  himself  with  the  character  of  the 
different  parts  of  the  flame,  and  the  use  of  the  blow-pipe  in  producing  the 
reducing  and  oxidizing  flames. 

§291.  The  flame  of  the  candle,  or  of  the  gas-jet,  burning  under  ordinary  circum- 
stances, consists  of  three  distinct  parts:  a  dark  nucleus  or  zone  in  the  centre, 
surrounding  the  wick,  consisting  of  unburnt  gas — a  luminous  cone  surrounding 
this  nucleus,  consisting  of  the  gases  in  a  state  of  incomplete  combustion.  Ex- 
terior to  this  is  a  thin,  non-luminous  envelope,  where,  with  a  full  supply  of 
oxygen,  complete  combustion  is  taking  place:  here  we  find  the  hottest  part  of 
the  flame.  The  non-luminous  or  outer  part  is  called  the  oxidizing  flame;  the 
luminous  part,  consisting  of  carbon  and  unconsumed  hydrocarbons,  is  called 
the  reducing1  flame. 

§292.  77* r  flame  produced  &//  ihc  Wow-pipe  (or  Bunsen  burner)  is  divided  into 
two  parts:  the  oxidizing  flame,  where  there  is  an  excess  of  oxygen,  correspond- 
ing to  the  outer  zone  of  the  candle-flame;  and  the  reducing  flame,  where  there 
is  an  excess  of  carbon,  corresponding  to  the  inner  zone  of  the  candle-flame. 
Upon  the  student's  skill  in  producing  these  flames  depend  very  largely  the 
results  in  the  use  of  the  blow-pipe. 

In  order  to  produce  a  good  oxidizing  flame,  the  jet  of  the  blow-pipe  is  placed 
just  within  the  flame,  and  a  moderate  blast  applied — the  air  being  thoroughly 
mixed  with  the  gas,  the  inner  blue  flame,  corresponding  to  the  exterior  part 

*  Jf  the  unknown  is  a  solution  the  solid  may  be  obtained  by  evaporation  on  the  water  bath. 


376  PRELIMINARY  EXAMINATION  OF   SOLIDS.  ^293. 

of  the  candle-flame,  is  produced:  the  hottest  and  most  effective  part  is  just 
before  the  apex  of  the  blue  cone,  where  combustion  is  most  complete. 

The  reducing1  flame  is  produced  by  placing-  the  blow-pipe  just  at  the  edge  of 
the  flame,  a  little  above  the  slit,  and  directing  the  blast  of  air  a  little  higher 
than  for  the  oxidizing-  flame.  The  flame  assumes  the  shape  of  a  luminous  cone, 
surrounded  by  a  pale-blue  mantle;  the  most  active  part  of  the  flame  is  some- 
what beyond  the  apex  of  the  luminous  cone. 

§293.  "The  blast  with  the  blow-pipe  is  not  produced  by  the  lungs,  but  by  1 1n- 
action of  the  muscles  of  the  cheek  alone.  In  order  to  obtain  a  better  knowledge 
of  the  management  of  the  flame,  and  to  practise  in  producing  a  good  reducing 
flame,  it  is  well  to  fuse  a  small  grain  of  metallic  tin  upon  charcoal,  and  raising 
to  a  high  heat  endeavor  to  prevent  its  oxidation,  and  keep  its  surface  bright; 
or  better,  perhaps,  to  dissolve  a  speck  of  manganese  dioxide  in  the  borax  bead 
on  platinum  wire — the  bead  becoming  amethyst-red  in  the  outer  flame  and 
colorless  in  the  reducing  flame.  The  beginner  should  work  only  with  sub- 
stances of  a  known  composition,  and  not  attempt  the  analysis  of  unknown 
complex  substances,  until  he  has  made  himself  perfectly  familiar  with  the 
reactions  of  at  least  the  more  frequently  occurring  elements. 

The  amount  of  substance  taken  for  analysis  should  not  be  too  large;  a 
quantity  of  about  the  bulk  of  a  mustard-seed  being,  in  most  cases,  quite 
sufficient. 

The  physical  properties  of  the  substance  under  examination  are  to  be  first 
noted;  such  as  color,  structure,  odor,  lustre,  density,  etc. 

Heat  in  Glass  Tube  Closed  at  One  End, 

§294.  The  substance,  in  fragments  or  in  the  form  of  a  powder,  is  introduced 
into  a  small  glass  tube,  sealed  at  one  end,  or  into  a  small  matrass,  and  heat 
applied  gently,  gradually  raising  it  to  redness,  if  necessary  with  the  aid  of  the 
blow-pipe.  When  the  substance  is  in  the  form  of  a  powder  it  is  more  easily 
introduced  into  the  tube  by  placing  the  powder  in  a  narrow  strip  of  paper, 
folded  lengthwise  in  the  shape  of  a  trough;  the  paper  is  now  inserted  into  the 
tube  held  horizontally,  the  whole  brought  to  a  vertical  position,  and  the  paper 
withdrawn;  in  this  way  the  powder  is  all  deposited  at  the  bottom  of  the  tube. 
By  this  treatment  in  the  glass  tube  we  are  first  to  notice  whether  the  sub- 
stance undergoes  a  change,  and  whether  this  change  occurs  with  or  without 
decomposition.  The  sublimates,  which  may  be  formed  in  the  upper  part  of  the 
tube,  are  especially  to  be  noted.  Escaping  gases  or  vapors  should  be  tested  as 
to  their  alkalinity  or  acidity,  by  small  strips  of  moist  red  and  blue  litmus 
paper  inserted  in  the  neck  of  the  tube. 

Heat  in  Glass  Tuhe  Open  at  Both  Ends. 

§295.  The  substance  is  inserted  into  a  glass  tube  from  two  to  three  inches 
long,  about  one  inch  from  the  end,  at  which  point  a  bend  is  sometimes  made; 
heat  is  applied  gently  at  first,  the  force  of  the  air-current  passing  through  the 
tube  being  regulated  by  inclining  the  tube  at  different  angles.  Many  sub- 
stances undergoing  no  change  in  the  closed  tube  absorb  oxygen  and  yield 
volatile  acids  or  metallic  oxides.  As  in  the  previous  case,  the  nature  of  the 
sublimate  and  the  odor  of  the  escaping  gas  are  particularly  to  be  noted.  The 
reactions  of  sulphur,  arsenic,  antimony  and  selenium  are  very  characteristic; 
these  metals,  if  present,  are  generally  easily  detected  in  this  way  (§69,  7). 

Heat  in  Blow-pipe  Flame  on  Charcoal. 

§296.  For  this  test,  a  well-burned  piece  of  charcoal  is  selected,  and  a  small 
cavity  made  in  the  side  of  the  coal;  a  small  fragment  of  the  substance  is  placed  in 
the  cavity,  and,  if  the  substance  be  a  powder,  it  may  be  moistened  with  a  drop  of 
water.  The  coal  is  inclined  at  an  angle  of  about  twenty-five  degrees  and  the  flame 
made  to  play  horizontally  upon  the  assay.  The  substance  is  first  heated  in  the 
oxidizing  flame  and  then  in  the  reducing  flame.  Any  metallic  beads  which  form 
are  allowed  to  cool  and  carefully  examined  for  malleability,  fusibility  and  color. 


§300  PRELIMINARY  EXAMINATION  OF  SOLIDS.  377 

Any  escaping  gases  are  to  be  tested  for  their  odor;  the  changes  of  color  which  the 
substance  undergoes,  and  the  nature  and  color  of  the  coating  which  may  form  near 
the  assay,  are  also  to  be  carefully  noted.  Some  substances,  as  lead,  may  be 
detected  at  once  by  the  nature  of  the  coating. 

Ignition  of  the    Substance   previously  Moistened  with    a  Drop  of  Cobalt 

Nitrate. 

§297.  This  test  may  be  effected  either  by  heating1  on  charcoal,  in  the  loop  of 
platinum  wire,  or  in  the  platinum-pointed  forceps.  A  portion  of  the  substance 
is  moistened  with  a  drop  of  the  reagent,  and  exposed  to  the  action  of  the  outer 
flame.  When  the  substance  is  in  fragments,  and  porous  enough  to  absorb  the 
cobalt  solution,  it  may  be  held  in  the  platinum-pointed  forceps  and  ignited. 
The  color  is  to  be  noted  after  fusion.  This  test  is  rather  limited;  aluminum, 
zinc  and  magnesium  giving  the  most  characteristic  reactions. 

Fusion  with  Sodium  Carbonate  on  Charcoal. 

§298.  The  powdered  substance  to  be  tested  is  mixed  with  sodium  carbonate, 
moistened  and  placed  in  the  cavity  of  the  coal.  Some  substances  form,  with 
sodium  carbonate  at  a  high  heat,  fusible  compounds;  others  infusible.  Many 
bodies,  as  silicates,  require  fusion  with  alkali  carbonate  before  they  can  be 
tested  in  the  wet  way.  Many  metallic  oxides  are  reduced  to  metal,  forming 
globules,  which  may  be  easily  detected. 

When  this  test  is  applied  for  the  detection  of  sulphates  and  sulphides,  the 
flame  of  the  alcohol  lamp  is  to  be  substituted  for  that  of  the  gas-flame,  as 
the  latter  generally  contains  sulphur  compounds. 

Examination  of  the  Color  which  may  be  imparted  to  the  Outer  Flame. 

§299.  In  this  way  many  substances  may  be  definitely  detected.  The  test  may 
be  applied  either  on  charcoal  or  on  the  loop  of  platinum  wire,  preferably  in  the 
latter  way.  When  the  substance  will  admit  a  small  fragment  is  placed  in  the 
loop  of  the  platinum  wire,  or  held  in  the  platinum-pointed  forceps,  and  the 
point  of  the  blue  flame  directed  upon  it.  If  the  substance  is  in  a  powder  it  may 
be  made  into  a  paste  with  a  drop  of  water,  and  placed  in  the  cavity  of  the 
charcoal,  the  flame  being  directed  horizontally  across  the  coal.  The  color 
\vhich  the  substance  imparts  to  the  outer  flame  in  either  case  is  noted.  In 
most  cases  the  flame  of  the  Bunsen  burner  alone  will  suffice;  the  substance 
being  heated  in  the  loop  of  platinum  wire,  which,  in  all  cases,  should  be  first 
dipped  in  hydrochloric  acid  and  ignited,  in  order  to  secure  against  the  presence 
of  foreign  substances.  Those  salts  which  are  more  volatile  at  the  temperature 
of  the  flame,  as  a  rule  give  the  most  intense  coloration.  When  two  or  more 
substances  are  found  together  it  is  sometimes  the  case  that  one  of  them  masks 
the  color  of  all  the  others;  the  bright  yellow  flame  of  sodium,  when  present  in 
exces's,  generally  veiling  the  flame  of  the  other  elements.  In  order  to  obviate 
this,  colored  media,  as  cobalt-blue  glass,  indigo  solution,  etc.,  are  interposed 
between  the  flame  and  the  eye  of  the  observer.  The  appearance  of  the  flame 
of  various  bodies,  when  viewed  through  these  media,  enables  us  often  to  detect 
very  small  quantities  of  them  in  the  presence  of  large  quantities  of  other 
substances. 

Treatment  of  the  Substance  with  Borax  and  Microcosmic  Salt. 

§300.  This  is  best  effected  in  the  loop  of  platinum  wire.  This  is  heated  and 
dipped  into  the  borax  or  microcosmic  salt  and  heated  to  a  colorless  bead;  a 
small  quantity  of  the  substance  under  examination  is  now  brought  in  contact 
with  the  hot  bead,  and  heated,  in  both  the  oxidizing  and  reducing  flames.  Any 
reaction  which  takes  place  during  the  heating  must  be  noticed;  most  of  the 
metallic  oxides  are  dissolved  in  the  bead,  and  form  a  colored  glass,  the  color 
of  which  is  to  be  observed,  both  while  hot  and  cold.  The  color  of  the  bead 
varies  in  intensity,  according  to  the  amount  of  the  substance  used;  a  very 


378  CONVERSION   OF  SOLIDS  INTO  LIQUIDS.  §301. 

small  quantity  will,  in  most  cases,  suffice.  Certain  bodies,  as  the  alkaline 
earths,  dissolve  in  borax,  forming-  beads  which,  up  to  a  certain  degree  of  satura- 
tion, are  clear.  When  these  beads  are  brought  into  the  reducing  flame,  and  an 
intermittent  blast  used,  they  become  opaque.  This  operation  is  called  flaming. 
As  reducing1  agents,  certain  metals  are  employed  in  the  bead  of  borax  or 
microcosmic  salt.  For  this  purpose  tin  is  generally  chosen,  lead  and  silver 
being  taken  in  some  cases.  These  metals  cannot  be  used  in  the  loop  of  plat- 
inum wire,  as  they  will  alloy  the  platinum.  The  beads  are  first  formed  in  the 
loop  of  wire;  then,  while  hot,  shaken  off  into  a  porcelain  dish,  several  being  so 
obtained.  A  number  of  these  are  now  taken  on  charcoal  and  fused  into  a  large 
bead,  which  is  charged  with  the  substance  to  be  tested,  and  then  with  the  tin 
or  other  metal.  For  this  purpose  tin  foil  (or  lead  foil)  is  previously  cut  in 
strips  half  an  inch  wide,  and  the  strips  rolled  into  rods.  The  end  of  the  rod 
is  touched  to  the  hot  bead  to  obtain  as  much  of  the  metal  as  required.  Lead 
may  be  added  as  precipitated  lead  ("  proof-lead  "),  and  silver  as  precipitated 
silver.  By  aid  of  tin  in  the  bead,  cuprous  oxide,  ferrous  oxide  and  metallic 
antimony  are  obtained  and  otjier  reductions  effected,  as  directed  in  §77,  7, 
and  elsewhere. 

CONVERSION  OF  SOLIDS  INTO  LIQUIDS. 

§301.  Before  the  fluid  reagents  can  be  applied,  solids  must  be  dissolved.  To 
obtain  a  complete  solution,  the  following  steps  must  be  observed: 

First.  The  solid  remaining  after  removal  of  organic  matter,  or  in  the  absence  of 
organic,  the  entire  solid,  is  treated  as  follows  unless  it  is  a  metal  or  alloy  (see 
§303).  The  solid,  reduced  to  a  fine  powder,  is  boiled  in  ten  times  its  quantity  of 
water.  Should  a  residue  remain,  it  is  allowed  to  subside,  and  the  clear  liquid 
poured  off  or  separated  by  filtration.  A  drop  or  two  evaporated  on  glass,  or  clean 
and  bright  platinum  foil,  will  give  a  residue,  if  any  portion  has  dissolved.  If  a  solu- 
tion is  obtained,  the  residue,  if  any,  is  exhausted,  and  well  washed  with  hot  water. 

Second.  The  residue,  insoluble  in  water,  is  digested  some  time  with  hot 
hydrochloric  acid.  (Observe  §305.)  The  solid,  if  any  remain,  is  separated  by 
-filtration  and  wrashed,  first  with  a  little  of  this  acid,  then  with  water.  The 
solution,  with  the  washings,  is  reserved. 

Third.  The  well-washed  residue  is  next  digested  with  hot  nitric  acid. 
Observe  if  there  are  vapors  of  nitrogen  oxides,  indicating  that  a  metal  or  other 
body  is  being  oxidized.  Observe  if  sulphur  separates.  If  any  residue  remains 
it  is  separated  by  filtration  and  washing,  first  with  a  little  acid,  then  with 
water,  and  the  solution  reserved. 

Sometimes  it  does  not  matter  which  acid  is  used  first.  But  if  a  first-group 
base  be  present,  HNO3  should  be  added  first,  for  HC1  would  form  an  insoluble 
chloride.  If  the  substance  contain  tin  (especially  an  alloy  of  tin)  HNO, 
would  form  insoluble  metastannic  acid,  H10Sn5015  ,  in  which  case  HC1  should 
be  used  first. 

Fourth.  Should  a  residue  remain  it  is  to  be  digested  with  nitrohydro chloric 
acid,  as  directed  for  the  other  solvents. 

The  acid  solutions  are  to  be  evaporated  nearly  to  dryness,  and  then  redis- 
solved  in  water,  acidulating,  if  necessary,  to  keep  the  substance  in  solution. 

Fifth.  Should  the  substance  under  examination  prove  insoluble  in  acids,  it 
is  likely  to  be  either  a  sulphate  (of  barium,  strontium  or  lead);  a  chloride,  or 
bromide,  of  silver  or  lead;  a  silicate  or  fluoride — perhaps  decomposed  by  sul- 
phuric acid — and  it  must  be  fused  with  a  fixed  alkali  carbonate,  when  the  con- 
stituents are  transposed  in  such  manner  as  to  render  them  soluble.  The 
water  solution  of  the  fused  mass  will  be  found  to  contain  the  acid;  the  residue, 
insoluble  in  water,  the  metal,  now  soluble  in  hydrochloric  or  nitric  acids 
(compare  §266,  7). 

If  more  than  one  solution  is  obtained,  by  the  several  trials  with  solvents, 
the  material  contains  more  than  one  compound,  and  the  solutions,  as  sepa- 
r-ited  by  filtration,  should  be  preserved  separately,  as  above  directed,  and 
analyzed  separately.  The  separate  results,  in  many  cases,  indicate  the  original 
combination  of  each  metal. 


§303.  TREATMENT   OF  A   METAL  OR   AN   ALLOY.  379 

TREATMENT  OF  A  METAL  OR  AN  ALLOY. 

§303.  As  most  metals  are  converted  into  soluble  compounds  by  nitric  acid,  this 
is  the  best  acid  to  use  for  the  solution  of  metals  and  alloys.  Of  the  common  metals, 
all  are  oxidized  and  all,  except  tin,  arsenic  and  antimony  are  converted  into  nitrates 
which  are  soluble  in  water.  Arsenic  is  converted  by  strong  nitric  acid  into  arsenic 
acid  which  is  readily  soluble  in  water.  Antimony  is  converted  into  the  pentoxide 
which  is  insoluble  in  water  or  nitric  acid  and  tin  is  converted  into  metastannic 
acid  which  is  also  insoluble  in  water  and  nitric  acid.  When  an  alloy  containing 
antimony  and,  especially,  tin  is  treated  with  nitric  acid,  considerable  portions  of 
the  other  metals  remain  in  the  insoluble  residue  with  the  tin  and  antimony.  The 
action  of  nitric  acid  on  these  metals  may  be  represented  as  follows : 

6Sb  +  10HN03  =  3Sb2O5  +  10NO  +  5H2O  . 
15Sn  +  2OHNO3  +  5H2O  =  3Hi0Sn5Oi5  +  20NO  . 
3Pb  +  8HN03  =  3Pb(N03)2  +  2ND  +  4H2O  . 

Bi  +  4HNO3  =  Bi(NO3)3  +  NO  +  2H2O  . 

Gold  and  platinum  are  not  attacked  by  nitric  acid.  Of  the  rare  elements,  titan- 
ium and  tungsten  are  converted  into  insoluble  hydroxides  like  tin,  tungsten  being 
converted  into  the  yellow  insoluble  tungstic  acid  (H2WO.i) .  Alloys  may  also  con- 
tain such  insoluble  constituents  as  ferro-chrome  or  ferrosilicon  which  separate  as  a 
black  residue. 

Carbon,  as  well  as  gold  and  platinum,  also  remains  as  a  black  insoluble  residue. 

Method  of  Procedure. 

About  one  gram  of  the  metal  or  alloy  in  the  form  of  sawings  or  drillings  is  placed 
in  a  beaker.  Add  10  cc.  concentrated  nitric  acid  and  5  cc.  water.  Cover  the 
beaker  with  a  watch  crystal  and  warm  on  the  water  bath  until  action  ceases.  If 
any  of  the  alloy  remains  unacted  upon  add  a  little  more  nitric  acid  and  water  and 
continue  the  heating.  When  the  action  is  complete  evaporate  to  dryness  on  the 
waterbath.  Add  5  cc.  concentrated  HNO3  and  20  cc.  water,  digest  on  the 
water  bath,  filter,  and  wash  the  residue.  The  solution  is  tested  for  the  metals  of  all 
the  groups. 

The  residue  is  digested  in  strong  hydrochloric  acid  and  if  necessary  in  aqua  regia. 
If  a  residue  still  remains  it  may  be  silicic,  titanic  or  tungstic  acid.  Transfer  to  a 

§latinum  crucible  and  add  a  few  cc.  of  hydrofluoric  acid  and  warm  under  a  hood, 
ilicic  acid  is  volatilized.  Add  a  few  drops  of  H2SO4  and  warm  until  the  hydro- 
fluoric acid  is  volatilized.  Titanium  and  tungsten  dissolve  on  diluting  with  water. 
If  a  considerable  residue  still  remains,  it  is  probably  metastannic  acid.  Filter  this 
and  wash  and  transfer  the  paper  with  the  precipitate  to  a  porcelain  crucible,  sup- 
ported on  a  pipestem  triangle,  and  burn  the  paper. 

Add  six  times  the  weight  of  the  precipitate  of  a  mixture  of  equal  parts  of  sulphur 
and  dry  sodium  carbonate.  Place  the  lid  on  the  crucible  and  heat  with  a  small 
flame  of  the  Bunsen  burner  until  the  sulphur  melts.  Continue  the  heating  until 
the  blue  flame  of  the  burning  sulphur  which  escapes  around  the  lid  of  the  crucible 
is  almost  gone.  Allow  the  crucible  to  cool  with  the  lid  on.  Place  the  crucible 
with  its  contents  in  a  beaker  and  dissolve  the  fused  mass  in  warm  water.  The  tin 
will  be  in  the  solution  as  sodium  thiostannate.  Any  insoluble  sulphides  of  copper, 
lead,  etc.,  are  filtered  off  and  the  solution  acidified  with  hydrochloric  acid.  The 
precipitated  sulphides  may  be  examined  for  tin  and  antimony  or  added  to  the 
remainder  of  the  sulphides  of  the  sub-group. 

If  considerable  tin  is  present  in  the  alloy,  it  may  be  more  advantageous  to  dis- 
solve it  by  the  following  procedure.  One  gram  of  the  alloy  which  has  been  cut 
into  small  shavings  with  a  clean  knife  or  otherwise  powdered  is  placed  in  a  beaker 
and  15  cc.  concentrated  hydrochloric  acid  added.  The  beaker  is  warmed  on  the 
water  bath,  one  drop  of  concentrated  nitric  acid  is  added  and  the  heating  con- 
tinued. When  the  action  again  ceases,  another  drop  of  nitric  acid  is  added.  This 
is  repeated  several  times  until  the  alloy  is  completely  dissolved.  By  this  process 
all  metals  in  the  alloy  are  converted  into  chlorides.  Some  of  these  chlorides,  espe- 
cially lead  chloride,  are  insoluble  in  hydrochloric  acid  and  remain  as  a  white  crys- 
talline deposit.  The  tin  is  converted  into  stannic  chloride  and  remains  in  solution 


380  TREATMENT  OF  A  METAL  OR  AN  ALLOY.  §303,  B. 

unless  too  much  nitric  acid  has  been  added,  which  converts  the  tin  into  the  insoluble 
metastannic  acid. 

The  solution  is  diluted  with  an  equal  volume  of  water  and  the  insoluble  chlorides 
of  the  first  group  filtered  off  and  analyzed  in  the  usual  manner.  The  remaining 
metals  are  tested  for  in  the  filtrate,  as  usual.  This  solution  may  be  turbid,  espe- 
cially when  cold,  due  to  the  action  of  water  on  SbCl3  .  H2S  is  passed  without  filter- 
ing as  the  precipitate  is  readily  converted  into  the  sulphide. 

B. — The  Residue  Insoluble  in  Nitric  Acid. 

This  may  contain  gold  and  platinum  in  their  metallic  forms,  and  tin  *  and 
antimony  *  in  the  form  of  metastannic  and  antimonic  acids.  The  separation 
of  the  two  former  from  the  two  latter  depends  upon  the  fact  that  the  meta- 
stannic and  antimonic  acids  are  soluble  in  hydrochloric  acid,  forming  SnCl4 
and  SbCl5  . 

Digest,  therefore,  the  well-washed  residue  in  concentrated  hydrochloric  acid 
at  a  boiling  temperature  for  from  5  to  10  minutes;  then  add  at  once  an  equal 
volume  of  water  (to  dissolve  the  stannic  chloride),  and  bring  to  the  boiling 
point. 

If  gold  or  platinum  existed  in  the  original  metal  or  alloy  it  will  now  be 
found  in  the  form  of  a  dark-browrn  or  black  powder  or  mass,  insoluble  in  the 
hydrochloric  acid.  If  such  a  residue  exists,  decant  tcJiile  Jwt,  again  add  hydro- 
chloric acid,  heat,  and  again  decant. 

The  Hydrochloric  Acid  Solution. 

This  solution  may  have  a  turbid  appearance,  especially  when  cold,  due  to  the 
action  of  the  water  upon  the  SbCl5;  but  without  filtering  proceed  with  the 
separation  and  detection  of  the  tin  and  antimony  by  the  usual  process. 

The  Dark-colored  Residue. 

Add,  after  washing,  two  volumes  of  hydrochloric  and  one  of  nitric  acid: 
evaporate  almost  or  quite  to  dryness,  dissolve  in  a  small  quantity  of  water 
(to  obtain  a  concentrated  solution),  and  divide  into  two  portions. 

The  gold  and  platinum  have  been  dissolved  by  the  aqua-regia  formed,  and 
now  exist  as  auric  and  platinic  chlorides. 

First  Portion — Test  for  Gold. 

Dilute  with  at  least  ten  times  its  bulk  of  water;  add  a  drop  or  two  of  a  mix- 
ture of  stannous  and  stannic  chlorides;  a  purple  or  bro\vnish-red  precipitate 
(or  coloration),  purple  of  Cassius,  constitutes  the  test  for  gold. 

A  convenient  way  of  preparing  this  mixture  of  stannous  and  stannic  chlorides 
is  to 

(a)  Add  a  few  drops  of  chlorine-water  to  a  solution  of  stannous  chloride;  or 

(6)  Add  to  a  small  quantity  of  stannous  chloride  enough  ferric  chloride  to 
produce  a  faint  coloration. 

Second  Portion — Test  for  Platinum. 

Add,  without  dilution,  an  equal  volume  of  a  strong  solution  of  ammonium 
chloride.  The  formation,  either  at  first  or  on  standing,  of  a  lemon-yellow 
crystalline  precipitate,  consisting  of  the  double  chloride  of  platinum  and 
ammonium,  (NH4Cl)2PtCl4  ,  constitutes  the  test  for  platinum. 

Addition  of  alcohol  favors  the  precipitation. 

If  the  proportion  of  platinum  is  very  small,  the  mixture,  after  ammonium 
chloride  has  been  added,  should  be  evaporated  to  dryness  on  a  water-bath  and 
the  residue  treated  with  dilute  alcohol.  The  ammonium  platinic  chloride 
remains  behind  as  a  yellow  crystalline  powder. 

*  Traces  may  sometimes  be  dissolved. 


§309.  SEPARATION    OF  ACIDS  FROM  BASES.  381 

SEPARATION  OF  THE  ACIDS  FROM  THE  BASES. 

§304.  The  preliminary  examination  of  the  solid  material  in  the  dry  way  will 
give  indications  drawing1  attention  to  certain  acids.  Solutions  can  be  evapo- 
rated to  obtain  a  residue  for  this  examination.  Thus,  detonation  (not  the 
decrepitation  caused  by  water  in  crystals)  indicates  chlorates,  nitrates,  bro- 
mates,  iodates.  Explosion  or  deflagration  will  occur  if  these,  or  other  oxygen- 
fiiniishing  salts — as  permanganates,  chromates — are  in  mixture  with  easily 
combustible  matter  (§273,  7).  Hypophosphites,  heated  alone,  deflagrate  in- 
tensely. A  brownish-yellow  capo-r  indicates  nitrates  or  nitrites  (§241,  7);  a 
grcc<n  flame,  berates  (§221,  7).  The  odor  of  Inirning  sulphur:  sulphides,  sulphites, 
tliiosulphates,  or  free  sulphur.  The  separation  of  carbon  black:  an  organic  acid. 
The  formation  of  a  silver  stain:  a  sulphur  compound  (§266,  7). 

§305.  When  dissolving  a  solid  by  acids  for  work  in  the  wet  way,  indications 
of  the  more  volatile  acids  will  be  obtained.  Sudden  effervescence:  a  carbonate 
(oxalate  or  cyanate,  §228,  6).  Greenish-yellow  vapors:  a  chlorate  (§272). 
Brownish-yellow,  chlornitrous  vapors  on  addition  of  hydrochloric  acid:  a  nitrate. 
The  characteristic  odors:  salts  of  hydrosulphuric  acid,  sulphurous  acid,  hydro- 
bromic  acid,  hydriodic  acid,  hj'drocyanic  acid,  acetic  acid.  The  separation  of 
sulpluir:  a  higher  sulphide,  etc.  It  will  be  remembered  that  chlorine  results 
from  action  of  manganese  dioxide,  and  numerous  oxidizing  agents,  upon 
hydrochloric  acid. 

§306.  If  the  material  is  in  solution,  the  bases  will  be  first  determined. 
(Certain  volatile  acids  will  be  detected  in  the  first-group  acidulation — by  indica- 
tions mentioned  in  the  preceding  paragraph.)  Now,  it  should  first  be  con- 
sidered, what  acids  can  be  present  in  solution  with  the  bases  -found?  Thus,  if 
barium  be  among  the  bases,  we  need  not  look  for  sulphuric  acid,  nor,  in  a 
solution  not  acid,  for  phosphoric  a<-i<I, 

§307.  As  a  general  rule,  the  non-alkali  metals  must  be  removed  from  a 
solution  before  testing  it  for  acids,  unless  it  can  be  clearly  seen  that  they  will 
not  interfere  with  the  tests  to  be  made. 

Metals  need  to  be  removed:  because,  firstly,  in  the  testing  for  acids  by  precipi- 
tation, a  precipitate  may  be  obtained  from  the  action  of  the  reagent  on  the 
base  of  the  solution  tested,  thus:  if  the  solution  contain  silver,  we  cannot  test 
it  for  sulphuric  acid  by  use  of  barium  chloride  (and  we  are  restricted  to  use 
of  barium  nitrate).  And,  secondly,  in  testing  for  acids  by  transposition  with  a 
stronger  acid — the  preliminary  examination  for  acids — certain  bases  do  not 
permit  transposition.  Thus,  chlorides,  etc.,  of  lead,  silver,  mercury,  tin  and 
antimoiw,  and  sulphide  of  arsenic,  are  not  transposed  by  sulphuric  acid,  or 
not  promptly. 

§308.  If  neither  arsenic  nor  antimony  is  among  the  bases,  they  may  all  be 
removed  by  boiling  with  slight  excess  of  sodium  or  potassium  carbonate,  and 
filtering.  Arsenic  and  antimony,  and  all  other  bases  of  the  second  group,  may 
be  removed  by  warming  with  hydrosulphuric  acid,  and  filtering.  When  the 
bases  are  removed  by  sodium  or  potassium  carbonate,  the  filtrate  must  be 
exactly  neutralized  by  nitric  acid,  with  the  expulsion  of  all  carbonic  acid  by  boil- 
ing. Then,  for  nitric  acid,  the  original  substance  may  be  tested. 

§309.  The  separation  of  phosphoric  acid  from  bases  is  a  part  of  the  work 
of  the  third  group  of  metals,  and  is  explained  in  §§152  and  153.  For  removal 
of  boric  acid,  see  §221;  oxalic  acid,  §151;  and  silicic  acid,  §249,  6  and  8. 

The  non-volatile  cyanogen  acids  can  be  separated  from  bases  by  digesting 
with  potassium  or  sodium  hydroxide  (not  too  strong,  §§231  and  232),  adding 
potassium  or  sodium  carbonate  and  digesting,  and  then  filtering.  The  residue 
is  examined  for  bases,  by  the  usual  systematic  process.  The  solution  will 
contain  the  alkali  salts  of  the  cyanogen  acids,  and  may  contain  metals  whose 
hydroxides  or  carbonates  are  soluble  in  fixed  alkali  hydroxides, 


382 


PRELIMINARY  EXAMINATION  OF  SOLIDS. 


§310. 


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Absence  of  volatile  bodies  (including  combined  wrater), 
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Organic  compounds  blacken  from  separation  of  carbon, 
Cu  and  Co  salts  blacken  at  high  heat. 

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Hg  (§58,  7),  flrray,  easily  rubbed  to  globules. 
HgCL  first  melts,  then  forms  white  crystalline  sublimat 
HgCl  ,  without  melting,  forms  a  sublimate,  yellow  while 
HgS  ,  a  black  sublimate,  turning  red  on  trituration. 
As  ,  steel-gray  sublimate;  garlic  odor. 
AsoO3  sublimes  in  white  octahedral  crystals,  does  not  i 
As2S3  ,  sublimate  nearly  black  while  hot,  reddish-yelloir  irl 
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PRELIMINARY  EXAMINATION  OF  SOLID8. 


383 


How  when  solidified. 

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NH4  salts,  those  not  decomposing,  white  sublimate  (§207,  7). 
FeCl3  slowly  sublimes  as  a  reddish-yellow  stain  (§126,  7). 
S  ,  free  or  by  reduction  of  sulphide,  gives  reddish-brown  drops, 
H2C204  ,  a  heavy  white  vapor  and  crystalline  sublimate. 
I  ,  a  violet  vapor  and  blue-black  sublimate. 

3  substance  evolves  d  gas  or  vapor: 
0  indicates  the  presence  of  a  nitrate,  chlorate,  bromate,  iodai 
piece  of  coal  placed  upon  the  assay  glows  upon  being  heated. 
H2S  ,  from  hydrated  sulphides  and  some  sulphites,  blackens 

by  its  odor. 
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bleaching  effect. 
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litmus. 

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Oxides  of  Nitrogen,  from  nitrates  or  nitrites,  reddish-brown,  acr 

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Oxides  are  obtained  from  metals,  except  from  Ag  ,  Au  and  J 
S  and  sulphides  yield  S02  .  Eecognized  by  its  odor  and  action 
As  yields  a  sublimate  of  As,O3  .  Garlic  odor. 

Sb  yields  a  sublimate  (white),  of  Sb203  and  Sb205  . 
Bi  ,  a  sublimate,  dark-brown  while  hot,  lemon-yellow  when  col 

Te,  gray  sublimate  of  tellurous  anhydride  (Te02). 
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Hg  ,  sublimate  of  metallic  mercury. 

e  substance  decrepitates: 
Crystals  as  NaCl.  (If  finely  pulverized,  the  decrepitation  is  a 

e  substance  deflagrates: 
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PRELIMINARY  EXAMINATION  OF  SOLIDS. 


§310. 


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ther  salts,  even  in  large  quantities,  do  not  interfer 

a,  green  glass,  appears  orange-yellow;  moistened  \ 
delicate  (§206,  7). 

io/e^:  K  and  most  of  its  salts,  except  borates,  ph 
violet  flame,  distinguished  in  presence  of  very  sm 
Excess  of  the  latter  prevents  the  reaction;  Li  als 

of  sodium,  the  potassium  flame  appears  reddish- 
glass  (§205,  7). 

e<Z:  Ca  and  its  compounds  produce  a  yellowish-red 
Sr  and  many  of  its  salts  yield  a  crimson  flame,  mi 
Li  and  its  salts  produce  a  carmine-red  flame  (§2 

reaction;  potassium  does  not. 
reen:  Yellowish-green,  Ba  and  most  of  its  salts,  j 

Emerald-green,  Cu  and  most  of  its  compounds. 
Bluish-green,  B203  . 

Yellowish-green,  B,03  ,  best  obtained  by  the  ad 

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SUBSTANCES   BEFORE   THE  BLOW-PIPfi. 


§311. 


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(Bor 


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GROUPING   OF   THE  METALS. 


387 


388 


TABLE  FOR  THE  SEPARATION  OF  THE  METAL*. 


Pb 

Hg' 
Ag 

As'" 

As* 

Sb* 

Sb'" 

Sniv 

Sn" 

AiT" 

ptiv 

Hg" 


|  Cu' 

1  Cu' 

• 

2  Cd 


|    A' 

CO       r*«T 


-3  Fe' 

|  Fe/ 

0 

fe  Co 


Ni 

Mn" 

Zn 
Ba 
Sr 
Ca 
Mg 

K 

Na 
NH4 


§313.     TABLE  FOR  REVIEW  OF 


PbCI2 

HgCI 
AgCI 


HgCI 


AgCI 


f  H2S04  =  PbSO*  White. 

J  H2S      =  PbS  Black. 

I  K2Cr04=  PbCr04  Yellow. 

I  Kl        =  Pbl2  Yellow. 

"I    o    fNH2HgCI  +  Hg   Black. 


j    S    [(NH3)3(AgCI)2  \  Add    HN03-|AgCI    White. 


'   As2Ss       ^ 

(NH4)4As2S5 

2 

As2Ss 

'o 

2 

f   H3As04 

i  •• 

As  S 

(NHJ,AsSA 

cd 

As^SK 

eg 

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1        1 

Sb2Ss 

5 

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2 

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S 

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1  1 

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AuCI3 

2!0* 

f 

PtS2 

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Solution. 

S    5 

PtS2 

PtCI4 

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fig 

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Pb(N03)2       ^    j    PbS04    Confirm  b. 

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§313.  TABLE  FOR  THE  SEPARATION  OF  THE  METALS.  389 

3  SEPARATIONS  OF  THE  METALS. 


tsH3 


I3  and  Sb 


tx8o§/ 


H3As03  (•  Remove  AgN03  with  CaCI2  and  add  H2S  -j  As2S3    Lemon  yeliow. 
SbAg3     j-  Dissolve  in  hot  HCI,  dilute,  filter  and  add  H2S  -|  Sb2S3   Orange. 


fig 


/SnCI2      J  Test  with  HgCI2.     -j  HgCI,  White ;  or  Hg  Gray. 

Sb       I=i2-f  SbCI6    l^w     -j  SbCI5  reject  or  test  in  Marsh  apparatus. 

Au          LlSIJ     AuCI        Iff®  (4ii      \     Dissolve  in  nitroliydro-j       AuCI3.NH4CI       Evaporate 
AUb'3      f  ~f  f  J  *"      I  cl'loric  acid,  evaporate  to  \  and  ignite  to  Au°,  YellOW. 
r,i  *  «-  °        r»A«  '   5-i      1  fdryness    with   excess  of 

V    Pt       J  o*|  I  PtCI4    J  «  g"-  I  Pt     i  NH4CI  and  digest  with  j    (NH4)2PtCI6   Ignite  to  Pt°, 

alcohol.  I  Gray, 

to  green-  )    Evaporate  to  dryness  with  excess  of  HN03.   Dissolve  residue  in  NH4OH  and  add  to  an  excess 

utionaCk  i  of  HCL    Test  tnis  solution  with  Na2HP04  ^Ammonium  phosphomolybdate,  Yellow, 
d  and  test  with  SnCI2  and  Cu  wire, 
mation  of  Pbl2  or  PbCr04. 
l)3   AddhotK2Sn02  j-Bi°    Black. 

tt)2.2NH4OH.2NH4N03^  Deep  blue  solution  evidence  of  copper. 

For  traces  add  HC2H302  and  test  with  K4Fe(CN)J  Cu2Fe(CN)6   Red-brown. 


H)22NH4OH.2NI!4N03I  Add  KCN  till  blue  color  disappears,  then  H2S-|  CdS    Lemon-yellow. 
HCI  and  precipitate  with  (NH4)2C03  j  AI(OH)3    White,  gelatinous. 

Acidify  with  HC2H302  and  add  Pb(C2H302)2-j  PbCr04    Lemon-yellow. 

I  Dissolve  in  HCI  and  add  KCNS  {  Fe(CNS>3    Blood  red. 

I  Test  original  solution  (acid)  with  KCNS  for  Fe'"  and  with  K3FefCN)e  for  Fe"-j  Fe3[Fe(CN)6]2   B]ue. 


fa. 

l_l> 
r  a. 
1  b 

Test  with  borax  bead.    Blue  bead. 
Add  NaHC03  and  H202,  Green  solution. 
Test  with  borax  bead.    Brown  bead. 
Heat  with  J  M!,ni^    '  add  Kl. 
Br  and  NaOH  i  N|(OH)3  f     Free  |  in  C32 

} 

Or  add 
nitroso-/3 
naphthol. 

'  Co—  Red 
precipitate. 

Ni  - 
Solution. 

\  Test  with 

1  AddNH4OH 

'     filter  and 

borax 
add  H 

bead. 
9S  -j  NiS. 

Black 

l)2j  Boil  with  Pb02and  HN03[HMn04   Purple. 


)a-<  Add  H2S  -ZnS    White. 


Dissolve  in  HCI  and  add  H2S04]BaS04  White. 


V 

a  n 

\ 

•g  i  ..  1.  Add  CaS04,  set  aside  ten  minutes  -}  SrS04,  White. 

|0a)a 

SrC03 

CO 

Sr(C2H302i2 

"£  o,fl       Moisten  SrS04  with  HCI  and  apply  flame  test. 

**•£ 

|5 

^0.2 
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I302U 

5    ^ 

CaC03 

Ca(C2H302)2  j 

S^          Filter  and  add    CaC  0     white,  soluble  in  HCI. 
(  Nn4)2u2U4 

H,PO.,,  White. 

-Apply  flame  test  using  cobalt  glass.    Violet. 
i — After  removal  of  Mg  apply  flame  test,  yellow. 


.     FIRST  TABLE. 


}314. 


fe 

H 
H 

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's  t 

£  t 

w_  CO 


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§314, 


AOTDS. 


TABLE. 


301 


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ACIDS.    FIRST  TABLE. 


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§314. 


ACIDS.     FIRST  TABLE. 


393 


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ACIDS.     FIRST  TABLE. 


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§314. 


ACIDS.     FIRST  TABLE. 


CD  «M  rc  • 

_r-  O  QJ  CO 

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"*  CD  c3  CQ 


^ 

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bo 


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r^l 
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I s 

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d  co 

2  ^ 


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^    ,0 


PH     CO 


S'S^'S'S 


5    o 


o    o 


be 


I 


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p,      r> 
3     9 


ntrated 


§ 

o 


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rS   ^ 


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396 


ACIDS.     FIRST   TABLE. 


§314. 


bJD 


O 

S    o 


C3       C3 

P-i 
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s  .-s 


^^  -°  s 

^31<§ 

1  §.-§  I 

a  *  .B"J 

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£   o   g   S 

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0     r&  & 

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2  ^ 

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"p  nd  -g 

""  'S  * 

OT  rj 


- 


ansposed  by  concentrated 
liberated.  Acidify  and 


bJD    bC  3 


coo 


o    a 

' 


,       CO 

§8 

o  coo 


o 

s  * 

C03^ 

of  c 


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1 1 

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111 

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p 

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§314. 
6 


ACID*.     FIRST   TABLE, 


39' 


1 

09 


O      <D 

S  -5 

^     G 


8 


c    s  ~    c 


II 


g  IJS 


ted  su 
ed. 


ng  with  concent 
sulphate  being 


jl  8  a 

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398 


AC7DS.     SECOND   TABLE. 


§315, 


w 
o 

s 

S 

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Q 

Q 
£ 
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PQ 
H 

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00 


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3  §  -o 

^•"2  P 


5  a  -a  -s  o 

fslll 

si!?* 
._-_. 


§316. 


ACIDS.     THIRD   TABLE. 


399 


CO 


bfl    bfl 


Cu  u 

S          'S 

03  -*-> 


be 


« 


- 


-a- 


1  -Mils  II 


. 

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fl  " 

0  r 


tift 


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S  >,  s 

I-0  3 

£  s  d 


III 

T3     0)       ra 

1rf2 

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£  °   g 

ii' 

^  ft  > 

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S  £   « 

3    P,    A 

1-; 

3  -0    2 

Sg  8 

|I| 
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o  .Z    c 


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£ 


400 


ACIDS.    FOURTH   TABLE. 


§317,   1. 


$317.  TABLE  FOR  IDENTIFICATION  AND  SEPARATION  OF  THE  COMMONLY 
OCCURRING  ANIONS  ( ACIDS).* 


CO. 
S02 
N20 

H2S 


HCN 
C2H40, 


1.  Boil  the  material  with  dilute  HN03 .     There  results: 

Effervescence;  turbidity  in  a  drop  of  lime-water. 
Effervescence,  penetrating  odor. 
Effervescence,  red-brown  fumes,  odor. 

Odor,  blackening  of  paper  moistened  with  lead  acetate,  separa- 
tion of  sulphur  in  the  solution. 

Odor  )  Often  masked  by  the  others;  see  special  tests 

Vinegar  odor  j       below. 

2.  Boil  with  concentrated  Na2C03  solution;  all  cathions  (bases)  except 
the  alkalis  are  precipitated  as  carbonates  or  hydroxides  and  removed  by 
filtration.     The  filtrate  contains  all  the  anions  (acids)  and  the  excess  of 
C03"  .     Acidulation  with  HN03  sets  free  C02 ,  and  Si02  is  precipitated; 
identified  in  the  microcosmic  salt  bead.     The  filtrate  is  made  ammoniacal. 

3.  Ca(N03)2  solution  precipitates: 

insoluble;  H2S04  liberates  HF . 
soluble,  reappearing  with  NH3; 
decolors  KMn04  solution, 
h  Fe"  +  Fe""  +  OH'  gives 
Prussian  blue  on  acidifying, 
with  K'  ions  in  concentrated  solution  po- 
tassium bitartrate  precipitated. 
In  the  filtrate  from  the  above, 
H2S  precipitates  As2S3  at  once  in  the  cold. 

In  the  filtrate  from  the  above, 
H2S  slowly  precipitates  from  hot  solution 
S2  +  As2S3 . 

In  the  filtrate  from  the  above, 
ammonium  molybdate  gives  yellow  pre- 
cipitate; or  Mg"  +  NH4*  +  OH'  gives 
MgNH4P04. 

4.  In  the  filtrate  from  3.  Ba(N03)2  precipitates: 

CrO/'(Cr207")  as  BaCr04 ,  yellow,  soluble  in  HC1;  the  yellow  color  of  the 
solution  becoming  green  on  boiling  with  alcohol. 


F 

as  CaF2       "J  insoluble        in 

V  in  acetic      dilute 

C204" 

as  CaC204    j      acid;         HC1 

CN' 

as  Ca(CN)2    v 

/heated 
Pruss 

C4H406" 

as  CaC4H406 

£ 

with  K' 

*3 

tassiu 

8 

In  th 

HAs03" 

as  CaHAs03 

~(-j 

H2S  pre< 

HAs04" 


as  CaHAs04 


HP04"     as  CaHP04 


*  From  Chem.  Prakt.    Abegg  and  Herz  (1900),  Breslau,  Page  113 ;  reviewed  by  Fresenius,  Z.t 
1900,  39,  566. 


by  OH'  /  coloration, 

gives  with       \   Prussian 


§318,   2.  NOTES   ON  THE  DETECTION  OF  ACIDS.  401 

S04"  as  BaS04  \  f  unchanged,  remains  insoluble  in 

HC1. 

SiF6"  as  BaSiF6  \  insoluble          on        1  gives    off    SiF4 ,    which    deposits 
in  HC1 ;      ignition  \       Si02  in  a  drop  of  water ;  the 
residue,    BaF2 ,    is    soluble    in 
HC1. 

5.  The  nitrate  from  4.  is  exactly  neutralized  with  HNO;!*;  Zn(N03)2 
then  precipitates : 
Fe(CN)/"  as  Zn3[Fe(CN)6]2  brownish-yellow        dissolved         (  brown 

Fe(CN)/'"  as  Zn2Fe(CN)6       white 

Fe'"  and  H*  (  blue. 

G.  A  few  drops  of  the  nitrate  from  5.  are  treated  with  as  little  Fe'"  as 
possible : 

Eed         f   Fe(CNS)3        )          on        J   permanent  red  color, 
coloration  1    Fe(C2H302)3   j     heating     1   precipitate  and  colorless  solution. 
In  the  absence  of  CNS'  another  drop  is  tested  with  Ag'  for  the  halogens; 
if  a  precipitate  results  or  if  CNS'  is  present,  one  part  of  the  solution  is 
treated  with  CS2  and  a  little  Cl-water: 

I'  violet  coloration,  disappears  with  )          ,    ni 

,.  ...        v  much  Cl-water. 

Br  brown  coloration,  does  not  disappear  with      j 

The  second  portion  is  evaporated  to  dryness  with  K,Cr207 ,  fused,  and 
the  mass  after  cooling  distilled  with  concentrated  H,S04;  appearance  of 
oily  brown  drops  of  Cr02Cl2 ,  forming  Cr04"  with  water:  Cl'  . 

7.  A  concentrated  water-extract  of  the  original  substance  is  treated 
with  concentrated  H2S04  and  solid  FeS04  or  Fe"  solution,  prepared  cold; 
a  brown  coloration  shows  the  presence  of  NO/. 

The  anions  mentioned  above  to  some  extent  exclude  one  another,  being 
unstable  when  together  in  solution  owing  to  their  power  of  mutual  oxida- 
tion and  reduction,  e.g.,  S03"  and  S";  SO/  and  NO/;  NO/  and  CN'; 
NO/  and  S";  NO/  and  I';  NO/  and  HAs03";  .S"  and  HAs04"  ,  etc.  It  is 
to  be  noticed  that  this  always  simplifies  the  analytical  procedure. 


§318.  NOTES  ON  THE  DETECTION  OF  ACIDS. 

1.  The  precipitation  of  tartrates  by  calcium  salts  is  incomplete;  from 
calcium  sulphate  solution  a  precipitate  forms  slowly  or  not  at  all.  Calcium 
tartrate  is  soluble  in  the  cold  in  a  solution  of  KOH  ,  precipitating-  gelatinous 
on  boiling',  again  soluble  on  cooling  (separation  from  citrate).  Calcium  tartrate 
is  soluble  in  acetic  acid  (separation  from  oxalate). 
2.  A  number  of  basic  carbonates  give  almost  no  effervescence  when  treated 

*In  the  original  German  text  it  is  directed  to  use  HCl  at  this  point. 


402  NOTES   (XV   THE   DETECTION  OF  ACIDS.  §318,   3. 

with  acids.  To  detect  the  presence  of  small  amounts  of  carbonate,  it  is  recommended 
to  place  the  dry  powder  in  a  test-tube  and  fill  about  three-fourths  full  of 
distilled  wrater.  Close  the  test-tube  with  a  two-holed  rubber  stopper  contain- 
ing- a  thistle  tube  reaching-  nearly  to  the  bottom  of  the  test  tube,  and  a 
delivery  tube  reaching-  just  through  the  stopper.  Add  dilute  sulphuric  acid 
and  warm  gently.  The  carbonate  is  decomposed,  driven  from  the  solution, 
and,  owing  to  the  limited  air  space,  readily  passes  through  the  delivery  tube 
into  the  solution  of  calcium  hydroxide. 

3.  With  the  generation  of  an  abundance  of  CO2  ,  the  precipitate  first  formed 
in   the    Ca(OH)2    is   redissolved    (solution    of    lime    in    spring   water).     Boiling 
drives   off  the   excess   of   C02    and   causes   the   reprecipitation    of   the   CaCOs  . 
Barium  hydroxide  may  be  used  instead  of  calcium  hydroxide. 

4.  If  compounds  have  been  strongly  ignited  previous  to  solution  for  analysis, 
oxalates  cannot  be  present. 

5.  In  Table  H    (§315),   if   strong  oxidizing  agents   are  present,   as   KC1O3  , 
K,Cr2O7  ,  KMn04  ,  etc.,  the  oxalic  acid  will  be  decomposed  on  warming  with 
hydrochloric  acid.     This  may  be  avoided  by   adding  calcium   chloride  to   the 
solution,  neutral  or  alkaline  with  ammonium  hydroxide.     The  oxalate  will  be 
precipitated   and   thus  separated   from  the   oxidizing  agents.     After  filtering, 
the   precipitate   is   digested  with   dilute   acetic    acid,   filtered    and    the   filtrate 
tested  for  phosphate  with  ammonium  molybdate.     The  residue  is  dissolved  in 
hydrochloric   acid,   filtered   if   necessary    (calcium   sulphate    does   not   dissolve 
readily),  and  the  filtrate  made  alkaline  with  ammonium  hydroxide.     The  pre- 
cipitate   thus    obtained    is    washed,    dissolved    in    nitric    acid    and    tested    with 
potassium  permanganate.     The   filtrate  from  the   solution   after  the   addition 
of  calcium  chloride  is  acidified  with  hydrochloric  acid,  heated  to  boiling  and 
tested  for  sulphate  by  the  addition  of  a  few  drops  of  barium  chloride  (§317). 

6.  In  Table  H,   if  sulphites  or  thiosulphates  are  present,   the  solution   in 
hydrochloric  acid  must  be  heated  sufficiently  to  drive  off  all  the  sulphurous 
anhydride,  or  reactions  for  oxalates  will  be  obtained,  due  to  the  sulphurous 
acid  alone.     If  there  be  any  doubt  as  to  the  complete  removal  of  the  sulphur- 
ous  anhydride,   the   gas   evolved   by   the   reaction   of   the   potassium   perman- 
ganate should  be  passed  into  a  solution  of  calcium  hydroxide.     A  precipitate  of 
calcium  carbonate  at  this  point  is  positive  evidence  of  the  previous  presence 
of  oxalic  acid  or  oxalates. 

7.  Alkali   ferro-   and   ferricyanides   are   separated   from   each   other  by   the 
solubility  of  the  latter  in  alcohol. 

8.  In   testing  for  nitric   acid   the   student   must  not   be   content   with   good 
results  from  one  test.     At  least  four  tests  should  be  made,  and  all  of  them 
should  give  positive  results  before  final  affirmative  judgment  is  passed.     Failure 
to  bleach  indigo  solution  in  the  presence  of  an  excess  of  hydrochloric  acid  may  be 
taken  as  conclusive  evidence  of  the  absence  of  nitrates. 

9.  In  the  analysis  of  minerals,  silica  or   silicates  will   usually  be   present. 
The  silica  should  be  removed  before  proceeding  with  the  analysis.     Fuse  the 
finely  divided  material  with  an  excess  of  sodium  carbonate,  digest  the  cooled 
mass  thoroughly  in  hot  water,  filter  and  evaporate  the  filtrate   to   dryness. 
Moisten  the  residue  with  concentrated  hydrochloric  acid,  and  again  evaporate 
to    dryness.     Pulverize    thoroughly,   digest    in    water    acidulated    with    hydro- 
chloric acid  and  filter.     The  residue,  white,  consists  of  the  silica,  SiO2  . 

10.  Meta-  or  pyrophospliates  do  not  react  promptly  with  ammonium  molyb- 
date.    In  the  usual  course  of  analysis  they  are  changed  to  the  orthophosphate 
(§255,  6A). 

11.  Phosphoric  acid  may  be  detected  in   the  presence  of  arsenic  acid  by 
ammonium  molybdate  if  the  solution  be  kept  cold;  it  is  preferable  to  remove 
the  arsenic  before  testing.     In  absence  of  interfering  substances  the  color  of 
the  silver  nitrate  precipitate  will  indicate  the  presence  or  absence  of  arsenic 
acid  (§69,  6;).     See  also  note  26. 

12.  Sulphides  which  are  transposed  bij  hydrochloric  acid  are  best  detected  by 
the  odor  of  the  evolved  gas,  and  by  passing  the  evolved  gas  into  ammonium 
hydroxide   and    testing  with    sodium   nitroferricyanide.     Other   sulphides   are 
decomposed  by   nitric  acid   or  by   nitrohydrochloric  acid   with   separation   of 
sulphur  as  a  leathery  mass  or  as  a  yellow  precipitate.     Persistent  heating  of 


§318,  15.  NOTES   ON   THE  DETECTION  Of  ACIDS.  403 

the  sulphur  with  the  reagent  decomposing"  the  sulphide  will  cause  the  oxida- 
tion of  a  portion  of  the  sulphur  to  a  sulphate  which  may  be  detected  in  the 
usual  manner.  A  portion  of  the  precipitated  sulphur  should  be  burned  on  a 
platinum  foil  with  careful  observance  of  the  odor  of  the  evolved  gas. 

1,}.  \  sulphite  (or  other  lower  oxidized  compound  of  sulphur)  is  readily 
detected  by  adding  barium  chloride  in  excess  to  a  portion  of  the  solution, 
dissolving  in  HC1  ,  filtering  if  residue  remains,  and  adding  bromine  or  chlorine 
to  the  clear  filtrate.  A  precipitation  of  barium  sulphate  indicates  the  oxidation 
of  a  lower  compound  of  sulphur  to  a  sulphate. 

I'l-  It  frequently  becomes  necessary  to  detect  and  estimate  sulphides,  thio- 
sulphates,  sulphites  and  sulphates  in  mixtures  containing  two  or  more  of  the 
compounds.  The  method  of  procedure  will  vary  according  to  the  nature  of 
the  substance.  The  student  will  be  aided  by  studying  §§257,  8;  258,  8;  and 
265,  8. 

15.  The  recognition  of  chlorides,  bromides  and  iodides — by  evolving  their 
chlorine,  bromine  and  iodine,  in  presence  of  each  other — can  be  accomplished  as 
follows — for  the  iodine  the  test  being  very  easy;  for  chlorine,  indirect  but 
unmistakable;  for  bromine,  dependent  upon  much  care  and  discretion.* 

The  iodine  is  liberated  with  dilute  chlorine-water,  added  drop  by  drop,  and 
is  readily  detected  by  starch,  or  carbon  disulphide,  according  to  §280,  8.  (As 
to  interference  of  thiocyanates,  see  §234.)  The  chlorine  is  vaporized  (from 
another  portion)  as  chlorochromic  anhydride,  and  the  latter  identified  by  its 
color  and  its  various  products,  as  described  in  §269,  Sd.  Before  the  bromine 
is  identified  the  iodine  is  to  be  either  removed  as  -free  iodine,  or  oxidized  to  iodate 
(§276,  86).  The  oxidation  to  iodic  acid  is  effected  as  follows:  Treat  with 
chlorine-water  till  free  iodine  no  longer  shows  its  color;  add  a  drop  or  two 
more  of  the  chlorine-water,  and  dilute  with  \vater,  keeping  cool;  then  add  the 
carbon  disulphide,  agitate  and  leave  the  solvent  to  settle,  for  the  yellow  color 
of  bromine.  The  removal  of  free  iodine  may  be  accomplished  as  follows:  Add 
chlorine-water,  drop  by  drop,  as  long  as  the  iodine  tint  seems  to  deepen  by 
five  addition;  add  the  carbon  disulphide,  agitate,  allow  to  subside,  and  remove 
the  lower  layer,  either  by  taking  it  out  with  a  pipette  or  by  filtration  through 
a  wet  filter.  Repeat,  if  need  be,  till  iodine  color  is  no  longer  obtained;  then 
continue,  with  dilute  chlorine  water,  in  test  for  bromine. 

Separation  by  the  persulphate  method.  To  ten  cc.  of  the  original  solution,  add 
slight  excess  of  Na2CO3  ,  free  from  chlorine,  and  boil,  to  precipitate  the  heavy 
metals.  The  solution  must  react  alkaline.  Filter  and  add  to  the  filtrate  acetic 
acid,  several  cc.  more  than  enough  to  neutralize  it,  dilute  to  50-60  cc.,  add  about 
one-half  gram  of  K2S2O8  ,  and  heat.  If  an  iodide  is  present,  free  iodine  will  be 
liberated,  and  may  be  identified  by  shaking  a  few  drops  of  the  solution  with  CF2  . 
Boil  in  a  casserole  until  all  iodine  is  expelled,  which  should  require  three  to  four 
minutes.  If  action  is  slow,  more  persulphate  should  be  added.  When  the  solu- 
tion is  colorless,  add  a  few  more  crystals  of  persulphate  and  boil  again,  to  make 
sure  that  no  iodine  remains.  As  the  solution  evaporates  add  distilled  water  to 
maintain  the  original  volume.  To  remove  Br'  add  two  cc.  of  F2SO4  ,  previously 
diluted  with  water,  a  little  more  K2S2O8 ,  and  heat  to  boiling  point,  but  do  not  boil. 
A  yellow  or  red  coloration,  if  the  separation  of  I  has  been  properly  conducted, 
indicates  Br  .  Pour  a  little  of  the  solution  into  a  test  tube,  cool,  and  shake  with 
CS2  ,  which  should  be  colored  yellow  or  red  but  not  violet,  which  would  indicate 
that  the  I  had  not  been  completely  removed.  If  bromine  is  present,  add  one-half 
gram  of  K2S2O8  to  the  main  part  of  the  solution,  and  boil  until  it  is  all  expelled 
and  the  solution  is  colorless;  then  test  with  a  little  more  K2S2O8  and  boil  five  min- 
utes longer  to  make  sure  of  the  complete  expulsion  of  the  bromine.  Be  sure  that 
the  volume  of  the  solution  does  not  fall  below  50  to  60  cc.  Add  distilled  water  from 
time  to  time  to  replace  that  lost  by  evaporation.  When  all  bromine  is  removed, 

*  In  consequence  of  the  relative  commercial  values  of  bromine  and  iodine,  and  the  medicinal 
relations  of  bromides  and  iodides,  it  is  of  great  importance  to  search  commercial  iodides  for 
intentional  and  considerable  mixtures  of  bromides— an  impurity  likely  to  escape  cursory 
chemical  examination.  There  are,  however,  very  slight  and  usually  unobjectionable  propor- 
tions of  bromides  frequently  to  b«  found  in  the  iodides  of  commerce,  and  occurring1  from  the 
difficulty  of  exact  s."pa-uti<  D  in  ;  !r-  ir.anu  hictiirc  of  k>«;ii!^  f  com  Ice! p. 


404  NOTES  ON   THE  DETECTION  OF  ACIDS.  §318,  16. 

cool  and  add  a  few  drops  of  silver  nitrate;  a  white,  curdy  precipitate  of  silver 
chloride  indicates  the  presence  of  Cl.  If  too  much  silver  nitrate  is  added,  a  white 
crystalline  precipitate  may  be  formed,  but  will  dissolve  upon  dilution  and  warming. 

If  ClO's  is  present,  the  above  procedure  cannot  be  followed,  for  the  I'  would 
be  oxidized  in  IO'3  .  In  this  case  it  is  necessary  to  precipitate  the  CF  ,  Br' ,  and  I' 
by  adding  to  the  original  solution  excess  of  silver  nitrate  and  then  nitric  acid; 
this  effects  a  separation,  silver  chlorate  being  soluble.  Wash  the  precipitate  of 
AgCl ,  AgBr  ,  Agl  ,  transfer  to  a  test  tube,  add  a  piece  of  zinc,  a  little  water,  and 
a  drop  of  sulphuric  acid  Let  it  stand  until  it  is  perfectly  black  all  the  way 
through,  showing  complete  reduction  to  metallic  silver.  Filter  and  treat  the 
filtrate  containing  ZnCl2 ,  ZnBr2 ,  ZnI2  ,  according  to  the  above  method,  starting 
at  the  beginning.  Even  if  no  heavy  metals  are  present,  Na^COs  should  be  added 
to  neutralize  any  mineral  acid  that  may  be  present  and  to  form  some  sodium 
acetate  when  acetic  acid  is  added. 

The  persulphate  method  should  be  used  only  when  the  presence  of  I'  or  Br' 
has  been  proved  by  some  short  test  (H2SO4,  Cl  ,  HNO2  ,  HNO3  ,  or  other  oxidizer). 
In  presence  of  a  great  excess  of  Br'  ,  CuSO4 ,  KNO2  ,  or  KgCl  is  an  excellent  test 
tor  I 

If  iodide  in  large  proportion  is  to  be  removed,  it  is  well,  first,  to  precipitate 
it  out,  as  far  as  possible,  by  copper  sulphate  and  a  reducing1  agent  (Note  11). 
The  filtrate  is  then  to  be  treated  by  either  method  above  given. 

If).  The  separation  by  ammonium  hydroxide,  as  a  solvent  of  the  silver  pre- 
cipitates— AgCl  ,  Ag'Br  ,  AgT — when  conducted  with  dilute  ammonium  hydrox- 
ide, may  be  made  complete  between  the  chloride  and  the  iodide,  also  between 
the  bromide  and  *the  iodide,  but  it  is  very  imperfect  between  the  bromide  and 
the  chloride.  The  hot  and  strong  solution  of  ammonium  acid  carbonate 
separates  the  chloride  from  the  bromide  (§269.  8r/). 

17.  The  direct  removal  of  iodides  In/  precipitation,  leaving  bromides  and  chlorides 
in  solution,  can  be  effected  (approximately)  by  copper  sulphate  with  sulphurous 
acid  (§77,  6f),  or  quite  completely  by  palladous  chloride  (§106,  6). 

78.  Chloric  acid  is  separated  from  hydrochloric  and-  all  other  acids  of  chlorine, 
bromine  and  iodine  (except  from  hypochlorous  acid,  and  from  traces  of  bromic 
acid),  bv  remaining  in  solution  during  the  precipitation  by  silver  nitrate 
(§273,  5). 

//).  Chloric  acid  is  separated  from  nitric  acid — after  finding  that  silver  nitrate 
(/Ires  no  precipitate  in  another  portion  of  the  solution,  acidulated — by  evaporat- 
ing and  igniting  the  residue,  then  dissolving  and  testing  one  portion  of  the 
solution  by  silver  nitrate  for  the  chloride  formed  from  chlorate  during  igni- 
tion (§273,  7).  The  other  portion  of  the  solution  is  tested  for  nitric  or  nitrous 
acid. 

20.  If  we  have  to  separate  chloric  acid  both  from  nitric  and  hydrochloric  acids, 
a  solution  of  silver  sulphate  must  be  used  instead  of  the  nitrate,  to  precipitate 
out  all  the  hydrochloric  acid.     The  filtrate  from  this  is   evaporated,  ignited, 
dissolved   and   tested   for   silver   chloride,    indicating   chlorate   in    the    original 
solution,   and   another  portion  is  tested   for  nitric   acid.     Also,   chlorates    are 
distinguished    (not  separated)    from  nitrates,  by   oxidation  of  ferrous   sulphate 
in  solution  with  acetic  acid  on  heating,  and  the  consequent  formation  of  the 
red  solution  of  ferric  acetate  (§§126,  6?>;  152;  223,  6).     The  solution  tested  must 
contain  110  free  acids,  and  no  nitrites  or  other'oxidizing  agents  beside  the  two 
in  question,  but  may  contain  chlorides;  and,  of  course,  the  ferrous  sulphate 
must   be  pure  enough  not  to  color  when  heated  alone  with  the   acetic  acid. 
Mix  the  ferrous  sulphate  solution  with  the  acetic  acid,  boil,  then  add  the  solu- 
tion to  be  tested,  and  heat  nearly   to  boiling,  for  some  minutes.     If  no  red 
color  appears,  chlorates  are  absent,  and  nitrates  may  be  present. 

21.  Hypochlorites  are  separated  with  chlorates  from  chlorides  (bromides),  etc., 
by  silver  nitrate;  and  distinguished  from  chlorates  (in  the  filtrate  from  AgGl  , 
etc.)   by  bleaching  litmus,  and  by  their  much  more  rapid  decomposition  and 
consequent  precipitation  of  any  silver  in  solution.     They  are  also  more  active 
than  chlorates,  as  oxidizing*  agents. 

8$.  M.  Dechan's  method  (§269,  8i)  consists  (/)  in  boiling  the  mixture  with  a 
solution  of  10  grammes  of  K2Cr2O7  ,  dissolved  in  100  cc.  of  water,  which  lib- 


§319,   1.  NOTES   ON    THE  DETECTION   OF   ACIDS.  405 

e rates  and  expels  all  oi  the  iodine  without  disturbing  the  bromine  and  chlo- 
rine 

5K2Cr20T  +  OKI  =  Cr203  +  8K2Cr04  +  3I2 

(2)  8  cc.  of  a  dilute  solution  of  sulphuric  acid  (consisting-  of  equal  volumes  of 
H^SO^  «;>.  (jr.  1.84,  and  water)  are  added  to  100  cc.  of  the  dichromate  solution, 
and  011  boiling,  the  bromine  is  distilled  off  without  disturbing-  the  chlorine; 
after  which  the  chlorine  is  detected  in  the  usual  manner.  For  other  methods 
of  detecting-  chlorides  in  presence  of  bromides  and  iodides,  see  §269,  8. 

23.  For  A.  Longi's  process  for  the  analysis  of  a  mixture  of  chlorides,  bro- 
mides,  iodides,   chlorates,   bromates,  iodates,   ferrocyanides   and   ferricyanides, 
see  C.  N.,  1883,  47,  209. 

24.  In  the  detection  of  chlorides  in  presence  of  cyanides  and  thiocyanates 
by   the   decomposition    of    the    silver   salts   with    concentrated    sulphuric    acid 
(§269,  8c),  a  drop  or  two  of  silver  nitrate  should  be  added  to  the  precipitate 
before  heating  with  the  acid  or  a  black  precipitate  will  be  obtained,  apparently 
carbon. 

2~).  For  the  detection  of  a  bromide  in  the  presence  of  an  iodide,  the  most 
satisfactory  method  is  by  the  use  of  potassium  chlorate  and  dilute  sulphuric 
acid  as  described  in  §276,  8c.  The  student  should  carefully  note  the  change 
in  color  of  the  solution.  The  first  very  dark  color  is  due  to  the  liberation  of 
iodine.  There  is  usually  a  sudden  change  of  color  on  the  complete  oxidation 
of  the  iodine,  but  if  much  bromine  be  present  the  solution  will  be  quite  dark 
straw  color.  This  should  be  tested  with  carbon  disulphide  and  the  heating 
continued  if  free  iodine  is  still  present. 

26.  Arsenic  acid  should  not  be  present  when  testing  for  a  phosphate.     If  the 
arsenic  acid  be  reduced  to  arsenous  acid  by  sulphurous  acid  it  will  not  interfere 
with  the  ammonium  molybdate  test  for  a  phosphate.     The  excess  of  sulphur- 
ous acid  must  be  removed  by  boiling  before  testing  for  the  phosphate.     Arsenic 
is  best  removed  by  precipitation  as  sulphide  in  the  second  group. 

27.  Chromic    acid    is    always    identified    by    reduction    and    precipitation    as 
chromic  hydroxide,  green,  in  the  third  group.     The  red  or  yellow  color  to  the 
original    substance    usually    gives    evidence    of    the   probable    presence    of    the 
hexad  chromium.     The  reduction  is  effected  in  the  course  of  analysis  by  hydro- 
sulphuric   acid   with   precipitation   of   sulphur.     It   is   advisable   to   reduce   all 
chromates  by  warming  with  hydrochloric  acid  and  alcohol  before  proceeding 
with  the  analysis.     Another  portion  of  the  substance  may  be  reduced  with 
sulphuric  acid  and  alcohol  and  tested  for  chlorides. 

28.  Manganates  are  readily  decomposed  by  water  with  formation  of  KMn04 
and  MnO,  .     In  the  presence  of  a  fixed  alkali  the  manganate  solution  is  green 
and  does  not  rapidly  change  to  permanganate.     The  mang-anates  and  perman- 
ganates in   solution  are  all  dark  colored    (green,  purple-red)    and   should   be 
reduced    by    warming    with    hydrochloric    acid    before    proceeding    with    the 
analysis. 

§319.  PEINCIPLES. 

In  the  following  statements,  the  term  salt  includes  only  cases  where 
the  metal  acts  as  a  base,  e.  g.,  chromium  salts  include  CrCl3 ,  not  K2Cr04 . 
Only  salts  of  ordinary  metals  are  included. 

1.  Hydroxides  when  brought  in  contact  with  acids  form  salts,  provided 
they  can  be  formed  by  any  means  in  the  presence  of  water.  The  same 
is  true  of  oxides.  But  A12S3  and  Cr2S3  are  not  formed  in  presence  of 
water.  (Some  oxides  after  ignition  fail  to  unite  with  all  acids,  e.  g.,  Sn02 , 
Fe203 ,  A1203 ,  but  by  long  boiling  unite  with  a  few  acids ;  while  ignited 
Cr203  is  insoluble  in  all  acids). 


406  PRINCIPLES.  §319,  'Z. 

2.  All  nitrates,  chlorates  and  acetates  are  soluble,  but  salts  of  cuprosum, 
bismuth,  tin,  antimony  and  the  oxysalts  of  mercury  require  some  free  acid 
to  hold  them  in  solution. 

3.  All  oxides  and  hydroxides  are  insoluble,  except  those  of  the  alkalis, 
those  of  Ba,  Sr  and  Ca  slightly  soluble.     The  fixed  alkalis  precipitate 
solutions  of  all  other  metallic  salts,  Ba,  Sr  and  Ca  incompletely.     The 
precipitate  with  silver,  antimonosum  and  mercury  is  an  oxide,  with  SnIV 
it  is  SnO(OH)2 ,  with  Sbv,  SbO(OH)3 ,  in  all  other  cases  a  normal  hydroxide. 
[At  boiling  heat  instead  of  normal  hydroxides  other  hydroxides  are  some- 
times formed,  e.  g.9  Fe4Os(OH)6 ,  and  Cu302(OH)2].     This  precipitate  re- 
dissolves   in   eight   cases,   forming,   if  potassium  hydroxide  be   used  .  .  . 
K2Pb02 ,  K2Sn02 ,  K2Sn03 ,  KSb02 ,  KSb03 ,  K2Zn02 ,  KA102 ,  KCr02 .     The 
last  precipitates  on  boiling. 

4.  Ammonium  hydroxide  precipitates  solutions  of  the  first  four  groups, 
manganese   and  magnesium   imperfectly   and   not   at   all   if   ammonium 
chloride  be  present.     The  precipitate  is  a  normal  hydroxide,  except  that 
with  SnIV  it  is  SnO(OH)2,  with  Sbv,  SbO(OH),,  with  Ag  and  Sb'"  the 
oxide,  with  Pb  a  basic  salt,  and  with  Hg  a  substituted  mercuric  ammonium 
compound,    Hg'  in  addition   forms  Hg° .     The  precipitate  redissolves  in 
six  cases,  viz.,   silver,   copper,   cadmium,   cobalt,   nickel  and  zinc.      Com- 
plex   ammonium    compounds    are   formed,    such    as    (NH3)2Agf  OH   and 
(NH,)4Zn++  (OH)  2. 

5.  The  chlorides  of  the  first  group  are  insoluble,  lead   chloride  slightly 
soluble.     Hydrochloric  acid  and  soluble  chlorides  precipitate  solutions  of 
salts  of  the  first  group,  lead  salts  incompletely,  and   normal  lead  salts  are 
not  precipitated  by  mercuric  chloride.     Cuprous  chloride  is  also  insoluble. 
(For  action  on  higher  oxides,  etc.,  see  §269,  QA.) 

6.  The  bromides  of  the  first  group  are  insoluble,  lead   bromide  slightly 
soluble  (less  than  the  chloride).  Hydrobromic  acid  and  soluble  bromides 
precipitate  solutions  of  the  salts  of  the  first  group,  lead  salts  incompletely. 
(For  action  on  higher  oxides,  etc.,  see  §276,  6.4.) 

7.  The  iodides  of  lead,   silver,   mercury  and   cuprosum    are    insoluble. 
Hydriodic  acid   and  soluble  iodides  precipitate  solutions  of  lead,  silver, 
ir.rrcury  and  cuprosum.     Cupric  salts  are  precipitated  as  cuprous  iodide 
witli  liberation  of  iodine.     Ferric   salts  are  reduced  to  ferrous,   arsenic 
arid  to  arsonous  acid,  Sbv  to  Sb'"  with  liberation  of  iodine. 

(Bismuth,  stannoiis  and  antimonous  iodides  are  really  insoluble  in  water,  and 
are  readily  formed  by  the  action  of  hydriodic  acid  or  soluble  iodides  on  the  dry 
or  merely  moistened  salts.  But  the  dissolved  salts  of  these  three  metals  fre- 
quently contain  so  much  free  acid  that  it  prevents  their  precipitation  by 
hydriodic  acid  or  by  soluble  iodides.  Arsenous  iodide  is  decomposed  by  water. 
It  may  be  formed  from  the  chloride,  best  by  adding  hydriodic  acid  or  a  soluble 
iodide  to  a  solution  of  arsenous  acid  in  strong-  hydrochloric  acid.  BisnuiUi 
iodide  is  black;  stunnous,  antimonous  and  arsenous  iodides  are  yellowish  reel..* 


§319,  1.1.  PRINCIPLES.  407 

8.  The  sulphates  of  lead,  mercurosum,  barium,  strontium  and  calcium  arc 
insoluble,  those  of  calcium  and  mercurosum  slightly  soluble.     Sulphuric 
acid   and   soluble   sulphates   precipitate   solutions    of   lead,   mercurosum, 
barium,  strontium  and  calcium;  calcium  and  mercurosum  incompletely. 

9.  (a)  The  sulphides  of  the  first  four  groups  are  insoluble.     Hydro- 
sulphuric  acid  transposes  salts  of  the  first  two  groups  in  acid,  neutral, 
and   alkaline   mixtures,   except   arsenic,   which   is   generally   imperfectly 
precipitated  unless  some  free  acid  or  salt  that  is  not  alkaline  to  litmus 
paper  be  present.     The  result  is  a  sulphide,  but  mercurosum  forms  mer- 
curic sulphide  and  mercury,  and  arsenic  acid  forms  arsenous  sulphide  and 
free  sulphur.     Ferric  solutions  are  reduced  to  ferrous  with  liberation  of 
sulphur.     In  acid  mixture  oth§r  third  and  fourth  group   salts  are  not 
disturbed,  but  from  solutions  of  their  normal  salts  traces  of  cobalt,  nickel, 
manganese,  and  zinc  are  precipitated.     (For  action  on  higher  oxides,  see 
§257,6.4). 

(b)  Soluble  sulphides  transpose  salts  of  the  first  four  groups.  The 
result  is  a  sulphide,  except  that  with  aluminum  and  chromium  salts  it  is 
a  hydroxide,  hydrosulphuric  acid  being  evolved.  With  mercurous  salts, 
mercuric  sulphide  and  mercury  are  formed;  with  ferric  salts,  ferrous 
sulphide  and  sulphur. 

10.  The  carbonates  of  the  alkalis  are  soluble.     Carbonates  of  the  fixed 
alkalis  precipitate  solutions  of  all  other  metallic  salts.     The  precipitate  is: 

a.  An  oxide  with  antimonous  salts. 

6.  A  normal  hydroxide  with  Sn",  Al ,  Cr'"  and  Fe'";  with  SnIV,  SnO(OH),  ; 
with  Sbv,  SbO(OH), . 

c.  A  normal  carbonate  with  Ba ,  Sr  and  Ca  salts  and,  if  cold,  with  silver, 
mercurosum,  cadmium,  ferrosum  and  manganosum. 

d.  A  basic  carbonate  in  other  cases,  except  mercuric  chloride,  which 
forms  an  oxychloride.     Carbonic  is  completely  displaced  by  strong  acids, 
for   example,   from   all   carbonates,   by   HC1 ,   HC103 ,   HBr ,   HBr03 ,   HI , 
HIOa,  H2C204,  KC2H302,  HN03,  H3P04 ,  H2S04 ,  and  even  by  H2S ,  com- 
pletely from  carbonates  of  first  four  groups,  incompletely  from  those  of  the 
fifth  and  sixth  groups  (Nandin  and  Montholon,  C.  r.f  1876,  83,  58). 

11.  All  normal  and  di-metallic  orthophosphates  are   insoluble   except 
those  of  the  alkalis.     The  normal  and  di-metallic  phosphates  of  the  alkalis 
precipitate  solutions  of  all  other  salts.     The  precipitate  is  a  normal,  di- 
metallic,  or  basic  phosphate,  except  that  with  mercuric  chloride  and  with 
the  chlorides  of  antimony  it  is  not  a  phosphate,  but  an  oxide,  or  an  oxy- 
chloride. 

All  phosphates  are  dissolved,  or  transposed  by  nitric,  hydrochloric  and 
sulphuric  acids,  and  all  are  dissolved  by  acetic  acid  except  lead,  aluminum 
and  ferric  phosphates.  All  are  soluble  in  phosphoric  acid  except  those  of 
lead,  tin,  mercury  and  bismuth. 


408  PRINCIPLES.  §319,  12. 

12.  Ignition. — a.  The  oxides  of  lead  and  iron  heated  in  the  air  to  a  red  heat 
form  Pb304  and  Fe203 ,  but  jat  a  white  heat  form  PbO  and  Fe304 .     Other 
oxides,  if  ignited  in  the  air  to  a  white  heat,  when  changed,  either  take  up 
or  lose  oxygen  and  leave  ultimately  the  following :  Ag ,  Hg ,  Au ,  Ft , 
Sn02,   Sb203,   As203,   Bi,03 ,   CuO ,   CdO ,  Fe304,     Cr203 ,    A1203,    CoO, 
NiO  ,  Mn304 ,  ZnO  ,  BaO  ,  SrO  ,  CaO  ,  MgO  ,  K20  ,  Na20  .     In  a  few  cases 
ignition  at  a  lower  temperature  gives  other  results,  e.  g.,  Co203 ,  Ba02 , 
Na202 ,  Sb204 ,  etc. 

&.  Alkali  hydroxides  ignited  in  air  at  a  white  heat  are  not  changed. 
Other  hydroxides  evolve  H20  and  leave  as  in  (a). 

c.  Alkali  carbonates  ignited  in  air  at  a  white  heat  are  not  changed. 
Other  carbonates  evolve  C02  and  leave  as  m  (a). 

d.  Fixed  alkali  oxalates  ignited  at  a  white  heat  in  absence  of  air  are 
changed  to  carbonates,  evolving  CO  .     Ba ,  Sr  and  Ca  oxalates  and  a  few 
others  at  a  red  heat,  in  absence  of  air,  form  carbonates  evolving  CO ,  at 
a  white  heat  these  carbonates  are  changed  to  oxides  evolving  C02 .     All 
oxalates  ignited  in  presence  of  air  at  a  white  heat  are  changed  as  in  (a), 
except  the  fixed  alkali  oxalates  which  are  left  as  carbonates.     In  all  cases 
when  air  is  present  the  CO  burns  to  C02 . 

e.  All  organic  salts  ignited  at  a  white  heat,  in  a  current  of  air,  leave 
residues  as  in  (a),  but  forming  carbonates  if  fixed  alkalis  are  present. 
The  products  evolved  depend  upon  the  composition  of  the  organic  por- 
tion of  the  salt. 

13.  The  following  acids  may  be  made  by  adding  sulphuric  acid  in 
excess  to  their  respective  salts  and  distilling: 

a.  Carbonic  from  all  carbonates, 
ft.  Nitric  from  all  nitrates. 

d.  Sulphurous  from  all  sulphites. 

e.  Hydrochloric  from  all  chlorides  except  those  of  mercury.     But  sul- 
phuric acid  transposes  the   chlorides  of  Ag,   Sn   and  Sb  with  extreme 
difficulty,  so  that  practically  other  methods  are  used  to  separate  hydro- 
chloric acid  from  the  chlorides  of  these  metals. 


§320. 


EQUATIONS. 


409 


§320.  EQUATIONS. 

It  is  recommended  that  in  the  class-room  some  attention  be  paid  to  the 
balancing  of  equations  as  representing  the  important  analytical  and  synthetic 
operations,  especially  those  involving  oxidation  and  reduction.  The  work  will 
be  simplified  by  a  careful  study  of  §§216,  217  and  218  and  application  of  the 
rule  as  illustrated  there.  When  the  time  permits,  the  operations  represented 
by  the  equations  studied  in  the  class-room  should  be  performed  by  each 
student  at  his  laboratory  work-table.  At  first  the  teacher  should  select  simpler 
equations  illustrating  analytical  operations  and  the  principles  (§319).  Later, 
the  more  difficult  equations  involving  oxidation  and  reduction  should  be  studied. 
The  student  should  give  the  authority  for  every  reaction.  The  accompanying 
list  of  equations  is  merely  suggestive  and  may  be  expanded  by  the  teacher  as 
the  time  permits.  In  each  equation  the  second  substance  is  to  be  considered 
as  in  excess;  that  is,  sufficient  to  produce  the  greatest  possible  change  in  the 
first  substance.  For  description  and  methods  of  making  the  basic  salts  used 
in  this  list,  see  Dammer's  Anorganishe  Chemie. 


1.  Sb  -f  HN03 

2.  As,  +  HN03 

3.  As,03  +  HN03 

4.  Mn(OH),  -f-  Pb02  +  HN03 

5.  MnS04  +  Pb304  +  H2S04 ,  dilute 

6.  MnOo  +  KN03  +  K2CO3  ,  fusion 

7.  Sa  +  KN03  +  K2C03  ,  fusion 

8.  MnS  +  KN03  +  K,C03  ,  fusion 

9.  Mn^  +  Pb304  -{-  HNO, 

10.  Cr(OH)3  +  KN03  +  K2C03  , 

fusion. 

11.  Pb3(As04)2  +  2n  4-  H2S04  ,  dilute 

12.  Cu2As2OT  -f  Zn  +  H2S04  ,  dilute 

13.  Pb(NO3)2  +  Al  +  KOH 

14.  Cu(NO3)2  +  Al  +  KOH 

15.  Bi(N03)3  +  Al  +  KOH 

16.  Hg1002(N03)8  +  Al  +  KOH 

17.  MnS  +  Mn(N03)2  +  K,C03  ,  fus. 

18.  Mn305  +  Pb304  +  HN03 

19.  Fe  +  HoS04  ,  con.,  hot. 

20.  KI  +  KI03  -f  H,S04  ,  dilute 

21.  MnS04  +  KMn04  +  H2S04  ,  dilute 

22.  (NaCl  +  K2CrO4  -f  H2S04),  dry, 

hot 

23.  KN03  +  FeS04  +  H2S04  ,  con., 

cold 

24.  K2Cr20(Cr04)3  +  HC1 ,  hot 

25.  Hg-80(N03)6  +  Al  +  KOH 

26.  Ag3As04  +  SnCl2  +  HC1 ,  sp.  gr. 

1.18 

27.  Pb02  +  K2C204  +  H2S04  ,  dilute 

28.  PbsO4  ,  white  heat 

29.  NaH,PO2  ,  ignition 

30.  Fe809(AsO3),  +  FeS  -f  HC1 

31.  FeBr2  +  HN03 

32.  Sn  +  HN03  ,  hot 

33.  KOH  +  Br2  ,  hot 

34.  FeI2  +  H2S04  ,  sp.  gr.  1.84,  hot 

35.  KBr  +  KBrO3  +  H2S04  ,  dilute 

36.  FeSO4  +  KMnO4  +  H2S04  ,  dilute 

37.  K2Cr20(CrOJ3  +  KOH  +  Br2 

38.  4Hg20,(N205)3  +  Al  -f  KOH 

39.  Ag-3AsO3  -f-  SnCl2  -f  HC1 ,  sp.  gr. 

1.18 

40.  Co203  ,  ignition,  white  heat 

41.  H2S  -f  HNOS  ,  sp.  gr.  1.42,  hot 


42.  Hg-3(As04)2  +  FeS  +  HC1 

43.  Fe8Ou(As03)2  +  KOH  +  Cla 

44.  FeI2  +  HN03  ,  sp.  gr.  1.48,  hot 

45.  Cr2(S04)3  -f  Cr(N03)3  +  K2C03  , 

fusion 

46.  Pb,(As04)2  +  Zn  +  H2S04  ,  dilute 

47.  KOH  +  C12  ,  cold 

48.  KBr  +  KI03  +  H2SO4  ,  dilute 

49.  (Cr2OHCl5  +  K2Cr04  +  H2SO4), 

dry,  hot 

50.  Zn403(N03)2  +  FeS04  +  H2S04  , 

concentrated,  cold 

51.  Hg3(As04)2  +  SnCL  +  HC1,  sp.gr. 

1.18 

52.  Mn3O6  ,  ignition 

53.  Fe202S08  +  Zn  +  H2SO4  ,  dilute 

54.  Bi2S3  +  HN03  ,  dilute,  hot 

55.  Hg3As04  +  FeS  +  HC1 

56.  Cr2(OH)4S04  +  KOH  +  C12 

57.  Fe(H2P02)2  +  HNO3 

58.  Cr203  +  KC103  +  K2CO3  ,  fusion 

59.  Cu502(As04)2  +  Zn  +  H2S04  ,  dil. 

60.  KOH  -f  C12  ,  hot 

61.  MnaOu  +  KC103  +  K2C08  ,  fusion 

62.  HIO,  +  SnCL  +  HC1 

63.  Bi12013(N03)10  +  FeS04  +  H2SO4  , 

con.,  cold 

64.  CrO3  ,  ignition 

65.  KMnO4  +  H2C204  +  H2S04,  dilute 

66.  FeAs04  +  SnCL  +  HC1 ,  sp.  gr.  1.18 

67.  Fe3Cl8  +  FeS  +  HC1 

68.  5CuO.As205  +  Fe  +  HC1 

69.  HIO3  -f  H2C204  ,  hot 

70.  (Cr2(OH)5Cl  +  K2Cr207  +  H,S04), 

dry,  hot 

71.  Fe(NO3)2  +  FeS04  +  H2SO4  ,  con., 

cold 

72.  Ag2SO4  +  Zn 

73.  H2S03  +  HN03 ,  sp.  gr.  1.42 

74.  FeAsO4  +  FeS  -f  HC1 

75.  Pb(AsO2)a  +  KOH  +  CL 

76.  Fe(NO3),  +  HNO3 

77.  MnsO5  +  Mn(NO,)2  +  K2COg  , 

fusion 

78.  Fe8O9(As03)2  -f  KOH  +  Br2 

79.  Pb1008(OH)8(NOs)6  +  Al  +  KOH 


410  PROBLEMS  IN  SYNTHESIS.  §321 

§321.  PROBLEMS  IN  SYNTHESIS. 

For  the  sake  of  a  more  thorough  drill  in  the  principles  of  oxidation  antf 
other  reactions,  a  few  problems  are  here  given;  a  part  of  them  the  student 
should  practically  work  at  his  table,  but  they  are  chiefly  designed  for  class 
exercises.  Special  care  should  be  taken  that  a  pure  product  be  formed  and 
that  the  ingredients  be  taken  from  the  sources  indicated.  In  each  case  the 
authority  for  every  step  in  the  process  should  be  stated. 

1.  Silver  oxide  from  metallic  silver. 

2.  Mercuric  bromide  from  mercurous  chloride  and  sodium  bromide. 

3.  Chromic  chloride  from  potassium  chromate  and  hydrochloric  acid. 

4.  Arsenic  acid  from  potassium  arsenite. 

5.  Potassium  arsenate  from  arsenous  oxide  and  potassium  hydroxide. 

6.  Lead  nitrate  from  lead  chloride  and  potassium  nitrate. 

7.  Mercurous  nitrate  from  mercuric  chloride  and  sodium  nitrate. 

8.  Mercurous  oxide  from  mercuric  oxide. 

9.  Mercuric  bromide  from  metallic  mercury  and  potassium  bromide. 

10.  Lead  nitrate  from  lead  dioxide  and  potassium  nitrate. 

11.  Lead  chromate  from  lead  hydroxide  and  chromium  hydroxide. 

12.  Barium  chromate  from  chrome  alum  and  barium  carbonate. 

13.  Mercuric  chromate  from  mercuric  sulphide  and  chromium  hydroxide. 

14.  Chromium  sulphate  from  potassium  dichromate  and  zinc  sulphide. 

15.  Phosphoric  acid  from  sodium  phosphate. 

16.  Phosphorus  from  calcium  phosphate. 

17.  Lead  iodate  from  sodium  iodide  and  lead  sulphide. 

18.  Silver  iodate  from  silver  chloride  and  iodine. 

19.  Ferric  arsenate  from  ferrous  sulphide  and  arsenous  oxide. 

20.  Mercuric  bromide  from  mercuric  sulphide  and  sodium  bromide. 

21.  Ammonium  sulphate  from  ammonium  chloride  and  sulphur. 

22.  Sodium  chloride  from  commercial  salt. 

23.  Phosphorus  from  sodium  phosphate. 

24.  Lead  sulphide  from  trilead  tetroxide  and  ferrous  sulphide. 

25.  Ferrous  sulphate  from  ferric  oxide  and  sulphuric  acid. 

26.  Ammonium  hydroxide  from  potassium  nitrate. 

27.  Cadmium  sulphate  from  cadmium  phosphate  and  ferrous  sulphide. 

28.  Mercurous  nitrate  from  mercuric  sulphide  and  nitric  acid. 

29.  Barium  sulphate  from  potassium  thiocyanate  and  barium  chloride. 

30.  Mercurous  chloride  from  mercuric  oxide  and  sodium  chloride. 

31.  Sodium  iodate  from  potassium  iodate  and  sodium  chloride. 

32.  Sodium  phosphate  from  calcium  phosphate  and  sodium  chloride. 

33.  Strontium  nitrate  from  sodium  nitrate  and  strontium  sulphate. 

34.  Potassium  sulphate  from  potassium  nitrate  and  sulphur. 

35.  Barium  sulphate  from  barium  chloride  and  zinc  sulphide. 

36.  Potassium  permanganate  from  manganese  dioxide  and  potassium  nitrate. 

37.  Arsenous  chloride  from  lead  arsenate  and  sodium  chloride. 

38.  Potassium  chromate  from  potassium  nitrate  and  lead  chromate. 

39.  Potassium  iodide  from  potassium  chloride  and  iodine. 

40.  Barium  chlorate  from  sodium  chloride  and  barium  nitrate. 

41.  Arsenous  sulphide  from  arsine  and  ferrous  sulphide. 

42.  Copper  sulphate  from  copper  sulphide. 

43.  Silver  nitrite  from  silver  chloride  and  sodium  nitrate. 

44.  Cuprous  chloride  from  metallic  copper  and  sodium  chloride. 

45.  Manganous  carbonate  from  manganese  peroxide  and  sodium  carbonate. 

46.  Manganous  pyrophosphate  from  manganese  peroxide  and  ammonium  phos- 

phate. 

47.  Lead  arsenate  from  lead  sulphide  and  arsenous  oxide. 

48.  Bismuth  subnitrate  from  metallic  bismuth  and  nitric  acid. 

49.  Barium  perchlorate  from  sodium  chloride  and  barium  hydroxide. 

50.  Lead  iodate  from  metallic  lead  and  iodine. 


§322.  TAPLE   OF   XOLrniLlTJES.  411 

§322.  TABLE  OF  SOLUBILITIES.* 

Showing  the  classes  to  which  the  compounds  of  the  commonly  occurriiu]  elements 

belong  in  respect  to  their  solubility  in  water,  hydrochloric  acid, 

nitric  acid,  or  aqua  regia. 

PRELIMINARY  REMARKS. 

For  the  sake  of  brevity,  the  classes  to  which  the  compounds  belong  are 
expressed  in  letters.  These  have  the  following  signification: 

W  or  w,  soluble  in  water. 

A  or  a,  insoluble  in  water,  but  soluble  in  hydrochloric  acid,  nitric  acid, 
or  in  aqua  regia. 

I  or  i,  insoluble  in  water,  hydrochloric  acid,  or  nitric  acid. 

Further,  substances  standing  on  the  border-lines  are  indicated  as  fol- 
lows: 

W — A  or  w — a,  difficultly  soluble  in  water,  but  soluble  in  hydrochloric 
acid  or  nitric  acid. 

W — I  or  w — i,  difficultly  soluble  in  water,  the  solubility  not  being 
greatly  increased  by  the  addition  of  acids. 

A — I  or  a — i,  insoluble  in  water,  difficultly  soluble  in  acids. 

If  the  behavior  of  a  compound  to  hydrochloric  and  nitric  acids  is  essen- 
tially different,  this  is  stated  in  the  notes. 

Capital  letters  indicate  common  substances  used  in  the  arts  and  in 
medicine,  while  the  small  letters  are  used  for  those  less  commonly  occur- 
ring. 

The  salts  are  generally  considered  as  normal,  but  basic  and  acid  salts, 
as  well  as  double  salts,  in  case  they  are  important  in  medicine  or  in  the 
arts,  are  referred  to  in  the  notes. 

The  small  numbers  in  the  table  refer  to  the  following  notes. 

Notes  to  Table  of  Solubilities. 

(1)  Potassium  dichromate,  W.  (2)  Potassium  borotartrate,  W.  (3)  Hydro- 
gen potassium  oxalate,  W.  (4)  Hydrogen  potassium  carbonate,  W.  (5)  Hydro- 
gen potassium  tartrate,  W.  (6)  Ammonium  potassium  tartrate,  W.  (7) 
Sodium  potassium  tartrate,  W.  (8)  Ammonium  sodium  phosphate,  W.  (9)  Acid 
sodium  borntp  W.  CIO)  Hydrogen  sodium  carbonate,  W.  (11)  Tricalcium 
phosphate,  A.  (12)  Ammonium  magnesium  phosphate,  A.  (13)  Potassium 
aluminum  sulphate,  W.  (14)  Ammonium  aluminum  sulphate,  W.  (15)  Potas- 
sium chromium  sulphate,  W.  (16)  Zinc  sulphide,  as  sphalerite,  soluble  in 
nitric  acid  with  separation  of  sulphur;  in  hydrochloric  acid,  only  upon  heating. 
(17)  Manganese  dioxide,  easily  soluble  in  hydrochloric  acid;  insoluble  in  nitric 
acid.  (18)  Nickel  sulphide  is  rather  easily  decomposed  by  nitric  acid:  very 
difficulty  by  hydrochloric  acid.  (19)  Cobalt  sulphide,  like  nickel  sulphide. 
(20)  Ammonium  ferrous  sulphate,  W.  (21)  Ammonium  ferric  chloride,  W. 

*  Tho  following:  table  of  solubilities,  is  taken  from  Fresenius  Qualitative  Analysis,  Well's 
translation  of  16th  German  edition. 


412 


TABLE  OF  SOLUBILITIES. 


SOLUBILITY 


I 
Potassium. 

Sodium. 

Ammonium. 

1 

3 

3 
1 

Strontium. 

Calcium. 

Magnesium. 

Aluminum. 

Chromium. 

d 

d 
N 

Manganese. 

Nickel. 

Cobalt. 

Oxide  

W 

W 

W 

W 

w 

W-A 

A 

A 

A&I 

A 

a,7 

A 

A 

Chromate. 

W, 

w 

w 

a 

w-a 

w-a 

w 

a 

w 

w 

a 

a 

Sulphate.. 

W18.,, 

W 

^14-2»'SO 

I 

I 

W-I 

W 

^13M4 

W&I16 

W 

W 

W 

W 

Phosphate 

W 

W8 

w8-n 

a 

a 

Au 

a,3 

a 

a 

a 

a 

a 

a 

Borate  .... 

Wj 

w» 

w 

a 

a 

a 

w-a 

a 

a 

a 

a 

a 

a 

Oxalate... 

W, 

W 

w 

a 

a 

A 

a 

a 

w-a 

a 

w-a 

a 

a 

Fluoride. 

W 

w 

w 

w-a 

w-a 

A-I 

a-i 

w 

w 

w-a 

a 

w-a 

w-a 

Carbonate 

W4 

W10 

W 

A 

A 

A 

A 

A 

A 

A 

A 

Silicate... 

W 

W 

a 

a 

a 

a 

a-i 

a 

a 

a 

a 

a 

Chloride.. 

w,7 

W86 

W,,.,. 

W 

W 

W 

W 

w 

W&I 

W 

W 

W 

W 

Bromide  .  . 

W 

w 

W 

w 

w 

w 

w 

w 

w&i 

w 

w 

AV 

w 

Iodide  

W 

w 

W 

W 

w 

w 

w 

w 

w 

w 

AV 

AV 

w 

Cyanide.  .  . 

W 

w 

w 

w-a 

w 

w 

w 

a 

A 

a 

a-i 

a-i 

FerrocyVle 

W 

w 

w 

w-a 

w 

w 

w 

A-I 

a 

i 

i 

Ferricy'de 

W 

w 

w 

w 

w 

a 

i 

i 

i 

Slphocy'de 

W 

w 

W 

-w 

w 

w 

w 

w 

w 

w 

w 

w 

Sulphide.  . 

W 

W 

W 

W 

w 

W-A45 

a 

a 

a-i 

A,. 

A 

a]8 

a,9 

Nitrate... 

W 

W 

W 

W 

W 

w 

w 

w 

W 

w 

w 

AV 

W 

Chlorate  .  . 

W 

w 

w 

w 

W 

w 

w 

w 

w 

w 

w 

W 

w 

Tartrate  .  . 

W»-«'T'1S«48 

W7 

w. 

a 

a. 

A 

w-a 

w 

w 

a 

w-a 

a 

w 

Citrate.... 

W 

w 

w 

a 

a 

w-a 

w 

w 

w 

w-a 

a 

w 

w 

Malate.... 

w 

w 

w 

w&a 

w 

w-a4T 

w 

w 

w 

w 

Succinate. 

w 

w 

w 

w-a 

w-a 

w-a 

w 

w-a 

w-a 

w 

AV 

w-« 

Benzoate.. 

•w 

w 

W 

w 

w 

w 

w 

Salicylate. 

ir 

W 

W 

w-a 

w-a 

w-a 

w 

Acetate... 

W 

W 

W 

W 

w 

W 

w 

W 

w 

W 

w 

W 

w 

Formate.. 

w 

w 

w 

w 

w 

w 

w 

w 

w 

w 

w 

W 

w 

Arsenite.. 

W 

w 

w 

a 

a 

a 

a 

a 

a 

a 

Arsenate.. 

W 

W 

w 

a 

a 

a 

a 

a 

a 

a 

a 

a 

a 

§322. 

TAFLE. 


TABLE  OF  SOLUBILITIES. 


413 


• 

d 

1  Ferrous. 

| 

I 

I 

Mercuroue 

Mercuric. 

•c 

a 

Bismuth. 

Cadmium. 

2 

Platinum. 

Stannous. 

Stannic. 

Antimonoi 

a 

a 

A34 

A 

A 

A 

a 

a 

a 

a 

a&i 

A42 

Qfcide 

w 

a 

A-I 

a 

w-a 

w 

a 

a 

a 

a 

Chromate 

W20 

V 

W-A 

A-I 

w-a 

W27 

W30 

w 

W 

w 

w 

a 

Sulphate 

a 

A 

a 

a 

a 

a 

a 

a 

a 

a 

a 

w-a 

Phosphate 

a 

a 

a 

a 

a 

a 

w-a 

a 

Borate 

a 

a 

a 

a 

a 

a 

a 

a 

w 

a 

w 

a 

Oxalate 

w-a 

v 

w 

a 

w-a 

a 

w 

w-a 

w 

w 

w 

Fluoride 

A 

a 

A 

a 

a 

A 

a 

a 

Carbonate 

a 

r 

a 

a 

a 

Silicate 

W 

"•  i 

I 

W-I 

A-I 

W28 

W 

W-A33 

W 

W85 

W37.38 

W 

W40 

W-A43 

Chloride 

w 

v. 

i 

w-i 

a-i 

w 

w 

w-a 

W 

w 

w 

w-a 

Bromide 

W 

i 

i 

W-A 

A 

A 

w 

a 

W 

a 

i 

w 

w 

w-a 

Iodide 

a-i 

I 

a 

W 

a 

a 

W 

w 

Cyanide 

i 

i 

i 

a 

i 

1 

i 

Ferrocy'de 

I 

- 

i 

w-a 

i 

Ferricy'de 

w 

w 

i 

a 

A 

w 

a 

w-a 

a 

w 

Sulphocy'de 

A 

a 

a23 

A 

A 

A2» 

a,i 

a 

A 

a38 

a39 

at, 

a4, 

•^44-45 

Sulphide 

w 

w 

W 

W 

wa« 

W 

W 

W* 

w 

w 

Nitrate 

w 

w 

w 

w 

w 

w 

w 

w 

w 

w 

Chlorate 

w-a 

W.j. 

a 

a 

w-a 

a 

w 

a 

w-a 

a 

*M 

Tartrate 

w 

vr 

a 

a 

a 

w-a 

w 

a 

Citrate 

w. 

w-a 

w-a 

a 

w-a 

w 

w 

w 

Malate 

w-a 

a 

a 

a 

a 

w-a 

w 

w 

a 

Succinate 

w 

R 

w-a 

a 

a 

w-a 

a 

w 

Benzoate 

w-a 

w-a 

w 

Saiicylate 

w 

V 

w 

W26 

w-a 

w 

W32 

w 

w 

w 

w 

Acetate 

w 

V 

w 

w-a 

w 

w 

w 

w 

w 

w 

Formate 

a 

il 

a 

a 

a 

a 

A 

a 

Arsenlte 

a 

B 

a 

a 

a 

a 

a 

a 

a 

a 

Arsenate 

TABLE  OF  SOLUBILITIES.  §322. 

(22)  Potassium  ferric  tartrate,  W.  (23)  Silver  sulphide,  only  soluble  in  nitric 
acid.  (24)  Minium  is  converted  by  hydrochloric  acid  into  lead  chloride;  by 
nitric  acid,  into  soluble  lead  nitrate  and  brown  lead  peroxide  which  is  insoluble 
in  nitric  acid.  (25)  Tribasic  lead  acetate,  W.  (26)  Mercurius  solubilis  Hahne- 
manni,  A.  (27)  Basic  mercuric  sulphate,  A.  (28)  Mercuric  chloride-amide,  A. 
(29)  Mercuric  sulphide,  not  soluble  in  hydrochloric  acid,  nor  in  nitric  acid,  but 
soluble  in  aqua  regia  upon  heating-.  (30)  Ammonium  cupric  sulphate,  W. 
(31)  Copper  sulphide  is  decomposed  with  difficulty  by  hydrochloric  acid,  but 
easily  by  nitric  acid.  (32)  Basic  cupric  acetate,  partially  soluble  in  water,  and 
completely  in  acids.  (33)  Basic  bismuth  chloride,  A.  (34)  Basic  bismuth 
nitrate,  A.  (35)  Sodium  auric  chloride,  W.  (36)  Gold  sulphide  is  not  dissolved 
by  hydrochloric  acid,  nor  by  nitric  acid,  but  it  is  dissolved  by  hot  aqua  regia. 
(37)  Potassium  plantinic  chloride,  W — I.  (38)  Ammonium  platinic  chloride, 
W — I.  (39)  Platinum  sulphide  is  not  attacked  by  hydrochloric  acid,  is  but 
slightly  attacked  by  boiling  nitric  acid  (if  it  has  been  precipitated  hot),  but 
is  dissolved  by  hot  aqua  regia.  (40)  Ammonium  stannic  chloride,  W.  (41) 
Stannous  sulphide  and  stannic  sulphide  are  decomposed  and  dissolved  by  hot 
hydrochloric  acid,  and  are  converted  by  nitric  acid  into  oxide  which  is  insoluble 
in  an  excess  of  nitric  acid.  Sublimed  stannic  sulphide  is  dissolved  only  by  hot 
aqua  regia.  (42)  Antimonous  oxide,  soluble  in  hydrochloric  acid,  not  in  nitric 
acid.  (43)  Basic  antimonous  chloride,  A.  (44)  Antimony  sulphide  is  com- 
pletely dissolved  by  hydrochloric  acid,  especially  upon  heating;  it  ?s  decom- 
posed by  nitric  acid,  but  dissolved  only  to  a  slight  degree.  (45)  Calcium 
antimony  sulphide,  W — A.  (46)  Potassium  antimony  tartrate,  W.  (47^  Hydro- 
gen calcium  malate,  W. 


§323.  REAGENTS.  415 

§323.  Reagents.* 

During  the  past  two  years  the  reagents  for  use  in  qualitative  chemical 
analysis  at  the  University  of  Michigan  have  been  made  up  on  the  basis 
of  the  normal  solution;  i.  e-.,  the  quantity  capable  of  combining  with  one 
gram,  of  hydrogen  or  with  its  equivalent  is  taken  in  a  litre  for  the  normal 
solution.  For  example:  Normal  potassium  hydroxide,  KOH  ,  requires  56.1 
grains  per  litre  of  solution  (not  56.1  grams  to  a  litre  of  water),  but  the  usual 
pure  product  contains  about  ten  per  cent  of  moisture,  so  it  is  directed  to 
use  62.3  grams  or  312  grams  for  a  solution  five  times  the  normal  strength, 
5K.  Barium  chloride,  BaCl2.2H20 ,  has  a  molecular  weight  of  244.2,  but 
the  hydrogen  equivalent  is  (244.2  -=-  2)  122.1,  so  for  a  litre  of  half-normal 
solution,  N/2,  take  61  grams. 

In  the  following  list  of  reagents,  in  the  parenthesis  immediately  follow- 
ing the  formula  are  given  the  grams  per  litre  necessary  for  a  solution 
of  the  strength  indicated.  Fresenius'  standard  follows  the  parenthesis. 

Acid,  Acetic,  HC2H8O2  (300,  5N),  sp.  gr.  1.04,  30  per  cent  acid. 

Arsenic,  H3AsO4.y0  H.O  (15,  ya  H3As04  ~  5). 

Flue-silicic,  HaSiF.  ,  §247. 

Hydrobromic,  HBr  (40,  N/2). 

Hydriodic,  HI  (64,  N/2). 

Hydrochloric,  HC1  (182,  5N,  sp.  gr.  1.084),  sp.  gr.  1.12,  24  p.  c.  acid. 

Hydrosulphuric,  H2S  ,  saturated  aqueous  solution,  §257,  4. 

lodic,  HIO3  (15,  i/2,  HI03-7-  6). 

Nitric,  HN03  (315,  5N,  sp.  gr.  1.165),  sp.  gr.  1.2,  32  p.  c.  acid. 

Nitrohydrochloric,  about  one  part  of  concentrated  HN08  to  three  parts 
HC1  . 

Nitrophenic,  C8H2(NO,)aOH  (picric  acid). 

Oxalic,  HoC2O4.2HoO  ,  crystals  dissolved  in  10  parts  water. 

Phosphoric,  H3PO4  (16,  N/2). 

Sulphuric,  H2SO4  ,  concentrated,  sp.  gr.  1.84. 

Sulphuric,  dilute  (245,  5N,  sp.  gr.  1.153),  one  part  acid  to  five  parts  water. 

Sulphurous,  H2SO3  ,  saturated  aqueous  solution. 

Tartaric,  H2C2H408  ,  crystals  dissolved  in  three  parts  water. 
Alcohol,  C2H60  ,  sp.  gr.  0.815,  about  95  p.  c. 
Aluminum  Chloride,  A1C1S  (22,  N/2). 

Nitrate,  Al(NO3)3.7y2H,O  (58,  N/2). 
Sulphate,  A12(SO4)3.18H2O  (55,  N/2). 
Ammonium  Carbonate,  (NH4)2CO8  (240,  5N),  one  part  crystallized  salt  in  four 

parts  water,  with  one  part  ammonium  hydroxide. 
Ammonium  Chloride,  NH4C1  (267,  5N),  one  part  salt  in  eig-ht  parts  water. 

Hydroxide,  NH4OH  (85NH8  ,  5N,  sp.  gr.  0.964),  sp.  gr.  0.96,  10  p.  c. 

NH3  . 
Ammonium  Molybdate,  (NH4)2MoO4  (36MoO,  ,  N/2,  §75,  6d),  150  g.  salt  in  one 

litre  of  NH4OH  ,  pour  this  into  one  litre  of  HN03  ,  sp.  gr.   1.2. 
Ammonium  Oxalate,  (NH4)2C2O4.i  H,<X  (40,  N/2),  one  part  crystallized  salt  in 

24  parts  water. 
Ammonium  Sulphate,  (NH4)2SO4  (33,  N/2). 

Sulphide,  (NH4)2S,  colorless,  three  parts  NH,OH ,  saturate  with 

H2S  and  add  two  parts  of  NH<OH  . 

*  In  the  greater  number  of  cases,  reagents  should  be  "chemically  pure."  Different  uses 
require  different  degrees  of  purity.  An  article  of  sodium  hydroxide  contaminated  with 
chloride  may  be  used  in  some  operations ;  not  in  others.  Those  who  have  had  training  in 
analysis  can  do  without  specific  directions,  which  cannot  be  made  to  cover  all  circumstances; 
and  the  beginner  must  depend  on  others  for  the  selection  of  reagents. 


REAGENTS.  §333. 

Ammonium  Sulphide,  (NH4)2SX  ,  yellow,  allow  the  colorless  to  stand  for  some 

time  or  add  sulphur. 
Antimonic  Chloride,  SbCl5  (30,  N/2). 
Antimonous  Chloride,  SbCl3   (38,  N/2). 

Arsenous  Oxide,  As2O3   (8,  N/4),  saturated  aqueous  solution. 
Barium  Carbonate,  BaC03  ,  freshly  precipitated. 

Chloride,  BaCL.2H2O  (61,  N/2),  one  part  salt  to  10  parts. water. 
Hydroxide,  Ba(OH)2.8H2O  (32,  N/5),  saturated  aqueous  solution. 
Nitrate,  Ba(NO3)o  (65,  N/2),  one  part  to  15  of  water. 
Bismuth  Chloride,  BiCl3  (52,  N/2,  use  HC1). 

Nitrate,  Bi(N03)3.5H,O  (40,  N/4,  use  HN03). 
Cadmium  Chloride,  CdCl2  (46,  N/2). 

Nitrate,  Cd{NO3)2.4H2O  (77,  N/2). 
Sulphate,  CdSO4.4H2O  (70,  N/2). 

Calcium  Chloride,  CaCl2.6H2O  (55,  N/2),  dissolve  in  5  parts  water. 
Hydroxide2  Ca(OH)2  ,  a  saturated  solution  in  water. 
Nitrate,  Ca(NO3),.4H2O  (59,  N/2). 

Sulphate,  CaSO4.2H2O  ,  a  saturated  solution  in  water. 
Carbon  Bisulphide,  CS2  ,  colorless. 
Chromic  Chloride,  CrCl3  (26,  N/2). 

Nitrate,  Cr(N03)3  (40,  N/2). 
Sulphate,  Cr2(SO4)3.18H2O  (60,  N/2). 
Cobaltous  Nitrate,  Co(NO3)2.6ELO  (73,  N/2),  in  8  parts  of  water. 

Sulphate,  CoS04.7H2O  (70,  N/2). 
Copper  Chloride,  CuCL>.2H,0  (43,  N/2). 

Nitrate,  Cu(NO3)2.6H2O  (74,  N/2). 
Sulphate,  CuSO4.5H2O  (62,  N/2),  in  10  parts  water. 
Cuprous  Chloride,  CuCl  (50,  N/2,  use  HC1). 
Ferric  Chloride,  FeCl3  (27,  N/2),  20  parts  water  to  one  part  metal. 

Nitrate,  Fe(NO3)3.9H2O  (67,  N/2). 

Ferrous  Sulphate,  FeS04.7H2O  (80,  N/2"),  use  a  few  drops  of  H2SO4  . 
Gold  Chloride,  HAuCl4.3H2O  ,  solution  in  10  parts  water. 
Hydrogen  Peroxide,  3  p.  c.  solution. 
Indigo  Solution,  6  parts  fuming-  H2SO4  to  one  part  indigo,  pulverize,  st'r  and 

cool,  allow  to  stand  48  hours  and  pour  into  20  parts  water. 
Lead  Acetate,  Pb(C2H302)2.3H2O  (95,  N/2),  dissolve  in  10  parts  of  water. 
Chloride,  PbCl,  ,  saturated  solution,  N/7. 
Nitrate,  Pb(NO3)2  (83,  N/2). 
Magnesia  Mixture:  MgSO4  ,  100  g-.;  NH4C1 ,  200  g-.;  NH4OH  ,  400  cc.;  H,O  ,  800 

cc.     One  cc.  =  0.01  g-.  P. 
Magnesium  Chloride,  MgCL.6H,O  (51,  N/2). 

Nitrate,  Mg(N03),.6H2O  (64,  N/2). 

Sulphate,  MgSO4.7H2O  (62,  N/2),  in  10  parts  of  water. 
Manganous  Chloride,  MnCl,.4Ho6  (50,  N/2). 

Nitrate,  Mn(N63),.6H,O  (72,  N/2). 
Sulphate,  MnSO4.7H2O  (69,  N/2). 

Mercuric  Chloride,  HgCl,  (68,  N/2),  in  16  parts  of  water. 
Nitrate,  Hg(N03),  (81,  N/2). 
Sulphate,  HgSO4  (74,  N/2). 
Mercurous  Nitrate,  HgN03   (131,  N/2),  one  part  salt,  20  parts  water  and  one 

part  HN03  . 
Nickel  Chloride,  NiCL.6H2O  (59,  N/2). 

Nitrate,  Ni(NO3)2.6H,O  (73,  N/2). 
Sulphate,  NiS04.6H2O  (66,  N/2). 

Palladous  Sodium  Chloride,  Na,>PdCl4  ,  in  12  parts  water. 
Potassium  Arsenate,  KsAs04  (26,  %  K3As04  -4-  5). 
Arsenite,  KAsO2  (24,  %~KAsOa  +  3). 
Bromate,  KBr03  (14,  i//KBr03  -4-  6). 
Bromide,  KBr  (60,  N/2). 
Carbonate,  KoCO3  (207,  3N). 
Chlorate,  KC10,  ,  the  dry  salt. 
Chloride,  KC1  (37,  N/2). 


§323.  REAGENTS.  417 

Potassium  Chromate,  K2CrO4  (49,  N/2),  in  10  parts  water. 
Cyanide,  KCN  (33,  N/2),  in  four  parts  water. 
Bichromate,  K1,Cr2O7  (38,  i/2,  K2Cr,O7  ^-4),  in  10  parts  water. 
Ferrocyanid'e,  k4Fe(CN)c.3H2O  (53,  N/2),  12  parts  water. 
Ferricyanide,  K3Fe(CN)6  (55,  N/2),  in  10  parts  water. 
Hydroxide,  KOH  (312  [90  p.  c.  KOH],  5N). 
lodate,  KIO3  (18,  %  KIO3  4-  6). 
Iodide,  KI  (83,  N/2),  dissolve  in  20  parts  water. 
Mercuric  Iodide,  K2HgI4  ,  Nessler's  solution,  §207,  6fc. 
Nitrate,  KNO.,  (50,  N/2),  the  crystallized  salt. 
Nitrite,  KNO,  ,  the  dry  salt. 
Pyroantimonate,  K2H2Sb2O7.6H20  ,  see  §70,  4c. 
Permanganate,  KMnO4  (16,  i/2  KMn04  -h  5). 
Thiocyanate,  KCNS  (49,  N/2),  in  10  parts  water. 
Hydrogen  Sulphate,  KHSO4  ,  fused  salt. 
Sulphate,  K,S04  (44,  N/2),  in  12  parts  of  water. 
Platinic  Chloride,  H,PtCl0.GH2O  ,  in  10  parts  of  water. 
Silve.  Nitrate,  AgNO3  (43,  N/4),  in  20  parts  of  water. 

Sulphate,  Ag,S04  ,  saturated  solution,  N/13. 
Sodiv.ni  Acetate,  NaC,H3O,,.3H2O  ,  in  10  parts  of  water. 

Carbonate,  Na2CO?    (159,  3N),  one  part  anhydrous  salt  or  2.7  parts  of 

the  crystals,  fra^COg.lOHoO  ,  in  5  parts  of  water. 
Chloride,  NaCl  (29,  N/2). 

Tetraborate,  Na2B4O7.10H2O,  lorax,  the  crystallized  salt. 
Hydroxide,  NaOH   (220   [90  p.  c.  NaOH],  5N),  dissolve  in  7  parts  of 

water. 

Hypochlorite  NaCIO,  §270,  4. 
Nitrate,  NaN03  (43,  N/2). 

Phosphate,  Na2HPO4.12H2O  (GO,  N/2),  dissolve  in  10  parts  of  water. 
Phosphomolybdate,  §75,  Qd. 
Sulphate,  (35,  N/2). 
Sulphide,  Na2S  ,  one  part  NaOH  saturated  with  HaS  to  one  part  of 

NaOH  ,  unchanged. 
Acid  Sulphite,  the  dry  salt. 

Sulphite,  Na2SO3.7H2O  (63,  N/2),  in  5  parts  of  water. 
Acid  Tartrate,  NaHC4H4O0  ,  in  10  parts  of  water 
Thiosulphate,  Na2S.,O.,.5H,O  ,  in  40  parts  of  water. 
Ststnric  Chloride,  SnCl4  (33,  N/2). 
Stanr.ous  Chloride,  SnCL.2HoO  (56,  N/2),  in  5  parts  water  strongly  acid  with 

HC1. 

Strontium  Chloride,  SrCL.GHoO  (67,  N/2). 
Nitrate,  Sr(NO3)2  "(53,  N/2). 
Sulphate,  SrS04  ,  a  saturated  aqueous  solution. 
Zinc  Chloride,  ZnCl2  (34,  N/2). 

Nitrate,  Zn(N03)2.6H,O  (74,  N/2). 
Sulphate,  ZnS04.7H2O  (72,  N/2). 


INDEX. 


Acetates,  detection  of 258 

ignition  of 267 

with  ferric  salts 157 

Acetic  acid 256-258 

estimation  of 258 

formation  of 257 

glacial 257 

occurrence  of 257 

preparation  of . 257 

properties  of 256 

reactions  of 257 

solubilities  of 257 

Acids,  detection  of,  notes  on 401 

displacement  of  weak  by  strong ...  185 
effect    of    concentrated    sulphuric 

upon 390 

list  of 13 

precipitated  by  barium  and  cal- 
cium chlorides 398 

preparation  of 408 

separation  from  bases 381 

table    of,    precipitated    by    silver 

nitrate 399 

table  of  separation  of 400 

Alabandite 177 

Alabaster 216 

Alkali   carbonates,   with   third   and 

fourth  group  salts 143 

group 227 

hydroxides,      action     on     double 

cyanides 272 

hydroxides,   detection  of  in  pres- 
ence of  carbonates 270 

hydroxides,  reactions  with 227 

Alkalis,   on  third  and  fourth  group 

metals 141 

Alkali  metals 5 

Alkaline  earth  metals 5 

earth  metals  in  presence  of  phos- 
phates    226 

earths,  relative  solubilities  of 210 

Alkali  sulphides,  as  reagents. . .  .317,  318 

action  of,  on  stannic  salts 86 

action  of,  on  stannous  salts 85 

Alloys,  analysis  of 379 

with  copper , , , ,  104 


Alluvial  sand 91 

Alpha  iron 154 

Aluminum 144-148 

acetate 146 

compounds,  ignition  of 147 

detection  of 147,  166 

distinction  from  chromium 150 

estimation  of 147 

hydroxide,    formation    and    prop- 
erties    145 

hydroxide,    solubility    in    ammo- 
nium chloride 164 

occurrence  of 144 

oxidation  of 148 

oxide  and  hydroxides 145 

phosphate,  separation  of 146,  147 

preparation  of 144 

properties  of 144 

reduction  of 147,  148 

salts,  reactions  of 145 

salts,  with  hydrosulphuric  acid. .  . .  146 

salts,  with  phenylhydrazin 146 

separation   of,   from    Cr   and   4th 

group  by  basic  acetates 145 

separation     of,     from     iron      by 

Na2S2O3  and  Na2SO3 146 

separation  of,  from  glucinum 201 

solubilities 144 

Alums 147 

Ammonia,  occurrence 235 

formation  of,  from  nitric  acid 286 

preparation  of 235 

properties  of 235 

Ammonium 235-239 

arsenomolybdate 62,  98 

benzoate,    in    separation    of    Cu 

from  Cd 107 

carbonate,  as  a  reagent 236 

carbonate,    in   separation    of   As, 

Sb  and  Sn 120 

chloride,  as  a  reagent 237 

chloride,  in  the  third  group 164 

chloride,  with  PtCl4 95 

compounds,  solubilities  of 235 

cyanate  in  formation  of  urea 279 

detection  of ,,,,,,,,,,,,,  , 238 


419 


420 


INDEX. 


Ammonium,  directions  for  detection  242 

estimation  of 238 

hydroxide,  as  a  reagent 236 

hydroxide,     as     a     distinguishing 

reagent  for  the  first  group 53 

hydroxide,  detection   by  mercuric 

chloride 238 

hydroxide,  preparation  and  prop- 
erties of 235 

molybdate,  preparation  of 98 

molybdate,  test  for  phosphates..  . .  311 

molybdate,  with  arsenic  acid 67 

oxidation  of 239 

phosphomolybdate 98 

picrate,  formation  of 236 

polysulphide,  formation  of 237 

salts,   detection   by   Nessler's   re- 
agent   237 

salts,  ignition  of 238 

solution  to  be  tested  for 242 

sulphate,   in  separation  of  stron- 
tium and  calcium 226 

sulphide,  as  a  reagent 237 

sulphide,  formation  of 236 

sulphide,  preparation  of 316 

sulphide,      on      iron      and      zinc 

groups 191 

sulphide,  yellow,  formation  of 115 

sulphide,  yellow,  in  separation  of 

cobalt  and  nickel 190 

sulphide,  yellow,  in  cupric  salts.  .  .    118 

test  for  nitric  acid 288-289 

thioacetate    as    a    substitute    for 

hydrosulphuric  acid 316 

Analysis  of  alkali  group 242 

proximate 14 

operations  of 13,  20 

ultimate 14 

Anatase 205 

Anglesite 29 

Anions,  table  of  separations  of 400 

Antimonic  acid 76 

distinction  from  antimonous 123 

reduction  to  antimonous  by  stan- 

nous  chloride 78 

salts,    action    of    hydriodic    acid 

on 78 

sulphide,  precipitation  of 77 

Antimonites 74 

Antimonous  argentide 79 

compounds  with  silver  nitrate ....     78 

iodide,  formation  of 78 

oxide,  formation  of 76 

salts  with  permanganates 78 

salts  with  chromates 78 


Antimonous  sulphide 74 

sulphide,  precipitation  of 77 

Antimony 72-82 

acids  of 72 

compounds,   reduction  with  char- 
coal       80 

detection  of,  in  alloys 379 

detection  of . 80 

detection  of  traces  of 123 

distinction  from  arsenic 78 

estimation  of 81 

in  the  test  for  aluminum 106-167 

metal  with  hydrosulphuric  acid ...     66 

mirror 65 

notes  on  analysis  of 123 

occurrence  of 72 

oxidation  of 81 

oxides  of 72 

pentachloride 74 

preparation  of 72 

properties  of 72 

reduction  of -. .     81 

reduction  to  metallic 79 

salts 74 

separation   from   arsenic   by  per- 
oxide of  hydrogen 121 

separation  from  arsenic 65 

separation    from    tin    by    sodiu.n 

thiosulphate 78 

separation  from  tin 81 

solubility  of 73 

spots 66 

sulphide,    separation   from   arse  i- 

ous  sulphide 123 

sulphide,    separation    from    sta  i- 

nous  sulphide 123 

with  iodine 66 

Apatite 216,  297 

Aragonite 216 

Argentite 46 

Argol,  purification  of 260 

Argyrodite 137 

Arrhenius 20 

Arsenates,   distinction  from  arsei- 

ites 70,  71 

separation  from  phosphates 402 

Arsenic 56-72 

acid,    precipitation    by   hydrosul- 
phuric acid 114 

acid,      reduction     by      hydrosi  1- 

phuric  acid  and  hydriodic  aci  i.     61 
acid,    reduction    with    sulphuro  is 

acid 60 

acid,  with  ammonium  molybdato.     67 
acid,  with  molybdates 62 


INDEX. 


421 


Arsenic  acid,  with  nitric  acid 66 

acid  with  silver  nitrate 67 

antidote  for 62 

compounds,  ignition  of 69 

compounds,      with      concentrated 

hydrochloric  acid 61 

compounds,  with  magnesium  salts  61 
compounds,  with    stannous    chlo- 
ride    89 

detection  of 70 

detection  of,  in  poisoning 68 

distinction  from  antimony 78 

estimation  of 70 

in  glass  tubing 70 

metal  with  hydrosulphuric  acid ...  66 

method  of  Fresenius  and  Babo. ...  68 

mirror 64,  65 

notes  on  analysis  of 122 

oxidation  of 71 

oxides  of 57 

occurrence  of 57 

pentasulphide,       formation      and 

properties  of *  60 

preparation  of 57 

properties  of 56 

reaction  with  alkali  sulphides 59 

reaction       with       hydrosulphuric 

acid 59 

reduction  of 71 

reduction  by  stannous  chloride. ...  61 

separation  from  antimony 65 

separation     from     antimony     by 

peroxide  of  hydrogen 121 

separation    from   Sb    and   Sn   by 

use  of  thiosulphates 60 

spots,  formation  of 64 

spots,  properties  of 66 

sulphide,      separation     of,     from 

Sb2S3 123 

sulphides     with    ammonium     car- 
bonate   120 

trichloride,   formation  in  analysis  61 

with  peroxide  of  hydrogen 71 

with  hydrosulphuric  acid  gas 67 

with  iodine 66 

with  nitric  acid 66 

Arsenites,    distinction    from    arsen- 

ates 123 

Arsenopyrite 57 

Arsenous  hydride 64 

oxide,  crystals,  identification  of .  .  .  67 

sulphide,  solubilities  of 58 

sulphide,  with  HC1  gas 67 

Arsine 64 

from  alkaline  mixtures 64 


Arsine  reactions  with  KOH 123 

separation  from  stibine 65 

with  hydrosulphuric  acid 60,  65 

Asbestos 299 

Asbolite 167 

Atomic  weights,  table  of 1 

Avogadro's  Hypothesis 21 

Azoimide  (hydronitric  acid) 282 

Azomide 8,  282 

Barite 211 

Barium 211-214 

carbonate,  action  on  ferric  salts.  156-157 

carbonate,  as  a  reagent 212 

carbonate,  as  a  reagent  for  third 

and  fourth  groups 143 

carbonate,  as  a  reagent  to  precip- 
itate chromium 150 

carbonate,  and  ferric  salts 156-7 

carbonate,      to      separate      phos- 
phates from  third,   fourth  and 

fifth  groups 194 

chloride,      separation      of,      from 

SrCl2  and  CaCl2  by  HC1 212 

detection  of 214 

estimation  of 214 

hydroxide,  formation  of 211 

iodide,  properties 370 

occurrence  of 211 

oxide,  preparation  of . .  211 

peroxide,  ignition  of 296 

peroxide,  preparation 211 

preparation  of 211 

properties  of 211 

salts,      separation      of      sulphites 

from  sulphates 213 

salts,  spectrum  of 213 

separation    of,    from  Sr,   Ca   and 

Mg  by  sulphates 213 

solubilities  of 212 

strontium    and    calcium,    separa- 
tion of  by  alcohol 226 

sulphate,  separation 215 

Bases,  alkali 12 

alkaline  earth 12 

copper,  group  of 11 

definition  of 3 

fifth  group  of 12 

first  group  of 11 

fourth  group  of 12 

iron  group  of 12 

need  for  separation  from  acids .  380,  381 

second  group  of 11 

silver  group  of 11 

sixth  group  of 12 


422 


INDEX. 


Bases,  third  group  of 12 

tin  group  of 11 

zinc  group  of 12 

Bauxite 144 

Beryl 200 

Beryllium 200 

Beta  iron 154 

Bismite 100 

Bismuth 100-104 

blowpipe,  reactions  of 103 

chloride,  sublimation  of. 103 

detection  as  iodide 103 

detection  by  alkaline  stanmte 103 

detection  by  cinchonine 102 

detection  in  alloys 379 

detection  of 103 

detection  of  traces  of 102 

dichromate 103 

estimation  of 103 

hydroxide,  solubility  in  glycerol ...    101 

iodide,  stability  toward  water 103 

nitrate,  precipitation  with  HC1. .  .  .    101 

nitrate,  reactions 101 

notes  on  analysis  of 130 

occurrence  of 100 

oxidation  of 104 

oxides  and  hydroxides  of 100 

oxychloride,  formation  of 101 

pentoxide,   reaction  with  halogen 

acids 101 

preparation  of 100 

properties  of 100 

reactions     of,     comparisons    with 

CuandCd 112 

reduction  by  grape  sugar 104 

salts,  reaction  with  the  alkalis ....    101 
separation  from  Cu  by  glycerol ...    101 

solubility  of 100 

sulphide,  formation  of 102 

sulphide,      separation     of,      from 

CuS 102 

sulphide,  separation  of,  from  tin 

group 102 

Bismuthinite 100 

Bismutite 100 

Bitter  spar 313 

Black  Band 155 

Blowpipe,  examination  of  solids. .  .  .   386 

Blue  vitriol 105 

Bonds,  plus  and  minus 244-245 

Borates,  green  flame  by  ignition  of  253 

in  analysis 54 

reactions  of 253 

Borax 252 

bead,  formation  of 254 


Borax,  bead,  test  for  Mn 189 

bead,  use  of 377 

Boric  acid 252-254 

estimation  of 254 

formation  of 252 

occurrence  of 252 

preparation  of 252 

properties  of 252 

solubility  of 253 

Boron 252 

Braunite 177 

Bromates,  detection  of 361 

estimation  of 362 

ignition  of 361 

preparation  of 360 

solubilities  of 361 

Bromic  acid 360-362 

properties  of 360 

reactions  of 361 

Bromides,  detection  of 359 

detection  in  presence  of  iodides. 

403,  404 

*  estimation  of 360 

formation  of 357 

ignition  of 359 

solubilities  of 357 

with  first  group  metals 358 

Bromine 354-356 

detection  of 356 

estimation  of 356 

formation  of 355 

occurrence  of 355 

preparation  of 355 

properties  of 354 

reactions  with 355 

solubilities  of 355 

Brookite 205 

Brown  ring,  test  for  nitric  acid 288 

Brucine,  reactions  with  nitric  acid .  .  290 

Brucite 220 

Cacodyl  oxide,  test  for  acetates  ....  258 

Cadmium 110-112 

detection  of 112 

estimation  of 112 

hydroxide 110 

notes  on  analysis  of 131 

occurrence  of 110 

oxide 110 

preparation  of 110 

properties  of 110 

reactions     of,     comparison     with 

Bi  and  Cu 112 

salts,   absorption  by  porous  sub- 
stances, separation  from  Cu.  ...  112 


INDEX, 


423 


Cadnuam  salts,  fused  with  K2S 112 

salts,  with  alkaline  tartrates,  sepa- 
ration from  Cu Ill 

salts,  with  alkalis Ill 

salts,  with  ammonia Ill 

salts,  with  barium  carbonate Ill 

salts,  with   pyrophosphates,  sepa- 
ration from  Cu Ill 

salts,     reactions      with     Na_>S2O;, 

separation  from  Cu 112 

salts,  reduction  of  by  metals 112 

salts,  reduction  of  by  ignition 112 

separation    from    Cu    by   H2S    in 

presence  of  KCN 107 

separation  from  Cu  by  KCNS Ill 

separation  from  Cu  by  glycerol ...    105 
separation   from    Cu   by   Na^S^O ; 

andNa2SO3 112 

solubilities  of 110 

Caesium 239-240 

Calamine 183 

Calaverite 138 

Calcium 216-220 

carbonate  in  spring  water 217 

carbonate,  solubility  of 224 

detection  of 219 

detection  of  by  spectrum 219 

estimation  of 219 

group. . 209 

group,  directions  for  analysis  of ...   224 
hydroxide,    formation    and    prop- 
erties    217 

hydroxide,  formation  by  Na2S.  ...   219 

hydroxide,  to  detect  CO2 218 

oxide,  formation  and  properties ...  216 

occurrence  of 216 

peroxide 217 

preparation  of 216 

properties  of 216 

salts,    separation    of    oxalic    from 

phosphoric  acid  by 218 

salts  with  Na2S 219 

separation    from    Ba   and    Sr   by 

amyl  alcohol 217 

separation    from    Ba   and    Sr   by 

(NH4)2SO4 217 

solubilities  of 217 

sulphate,    separation    from    stron- 
tium sulphate 215 

sulphate,      solubility     in     ammo- 
nium sulphate 226 

sulphate,  to  detect  strontium 219 

Calomel 37 

Carbon 254-256 

amorphous 255 


Carbon,  detection  of 256 

dioxide 267-271 

dioxide,  absorption  by  Ca(OH)2 .  . .  269 
dioxide,  detection  in  sodium  car- 
bonate   270 

dioxide,  detection  by  calcium  hy- 
droxide     218 

dioxide,     distinction     from      H2S, 

SO2,  N2O3,  etc 269 

dioxide,  formation  of 267 

dioxide,  occurrence  of 267 

dioxide,  properties  of 267 

monoxide 262-263 

preparation  of 255 

properties  of 254 

reactions  of 255 

reduction  by  ignition  with 256 

relations  of 10 

solubilities  of 255 

Carbonates,  acid,  decomposition  of  236 

decomposition  of,  by  acids 270 

detection  of 270 

detection  of  traces 402 

estimation  of 271 

ignition  of 270 

occurrence  of 267 

preparation  of 267 

reactions  with 268 

Carbonic  oxide,  formic  anhydride.  .  262 

Carnallite 228 

Cassiterite 82 

Cassius'  purple 93 

Castor  &  Pollux 239 

Celestite 214,  331 

Cement 299 

Cementite 154 

Cerargyrite 46 

Cerite 198,  199,  202 

Cerium 198 

Cerussite 29 

Chalcocite 104 

Chalcopyrite 45,  104 

Chalk 216 

Chamber  process  for  sulphuric  acid 

manufacture 331 

Chili  saltpeter 233,  363,  369 

occurrence  of 285 

Chloric  acid 350-353 

formation  of 350 

preparation  of 351 

properties  of 350 

separation  of,  from  nitric  acid ....   403 

Chlorates,  detection  of 353 

distinction  from  nitrates 404 

estimation  of .  353 


424 


INDEX. 


Chlorates,  formation  from  chlorine.   340 

ignition  of 352 

oxidation  by  ignition  of 352 

preparation  of 351 

reactions  with 351 

solubilities  of 351 

Chlorides,  detection  of 151 

detection  of,  in  presence  of  bro- 
mides  346,  347,  403 

detection   of,   in   presence   of   cy- 
anides or  thiocyanates 346,  405 

formation  of 341-342 

ignition  of 345 

Chloride  of  lime,  formation  of 348 

estimation  of,  by  H2O2 296 

Chlorine 337-341 

action  on  metals 339 

as  an  oxidizer 339 

detection  of 341 

estimation  of 341 

formation  of 338 

occurrence  of 338 

peroxide,    formation    and    proper- 
ties    350 

properties  of 337 

solubilities  of 338 

Chlorochromic  test  for  chlorides . . .  346 
anhydride 151 

Chlorous   acid,   formation  and   de- 
tection   349 

properties  of 349 

Chromates 151,  152 

in  test  for  HC1 151 

reduction     of,     by     hydrochloric 

acid 151 

reduction  of  by  H2S 150 

use  in  separation  of  barium 213 

with  antimonous  salts 78 

with  As'" 151 

with  ferrous  salts 161 

Chrome-ironstone 148 

Chromic  acid,  detection  of 152 

formation  of 151 

identification  of 405 

Chromite 148 

Chromium 148-153 

distinction  from  aluminum 150 

estimation  of 152 

hydroxide,  solubility  in  ammonium 

hydroxide 165 

and    manganese    in    third    group 

separation 166 

metal,  solubility  of 149 

occurrence  of 148 

oxidation  of .   152 


Chromium  oxides  and  hydroxides . .  148 

oxide,  solubilities  of 149 

preparation  of 148 

properties  of 148 

reduction  of 152 

salts,  solubilities  of 149 

salts,  reaction  of 149 

separation    from    Al    and    Fe    by 

H2O2 152 

separation  from  fourth  group 150 

separation    from    Fe   by   Na-jSgOs 

and  Na2SO3 146-147 

Chromous  salts 149 

Cinchonine  as  a  test  for  bismuth.  .  .   102 

Citric  acid 258-259 

detection  of  oxalic  acid  in 259 

distinction  from  tartaric 259 

properties  and  reactions 259 

Cinnabar 37,  313 

Clay  iron  stone 155 

Coal,  anthracite 255 

Cobalt 167-172 

bead  test 172 

cobalticyanide     separation     from 

nickel 169 

detection  of 172 

detection    of   by    means    of    am- 
monium thiocyanate 170 

detection    of   in    presence    of   Ni 

by  H2O2 190 

estimation  of 172 

hydroxide 165 

metal,  solubilities  of 168 

nitrate,  effect  of  ignition  with 377 

occurrence  of 167 

oxidation  of 173 

oxides  and  hydroxides 167 

phosphate,     a     distinction     from 

Ni 171 

preparation  of 167 

properties  of 167 

reduction  of 173 

salts,  solubilities  of 168 

salts,  with  alkalis 168 

salts,  with  barium  carbonate 169 

separation  from  nickel  by  ether ...  168 
separation  from  nickel  by  KNO2  170 
separation  from  nickel  by  KMnO4 .  172 
separation  from  nickel  by  NH4CNS  172 
separation  from  nickel  by  ni- 

troso-/3-naphthol 170,  190 

Cobaltite 57,  167 

Colloidal  sulphides  of  fourth  group  189 

Color,  flame  tests 377 

Columbite 198,203 


INDEX. 


425 


Columbium 198-199 

distinction  from  Ti 207 

properties  and  reactions  of .  .  .  .198-199 

separation  from  tantalum 203 

Contact  process  for  sulphuric  acid 

manufacture 331 

Copper 104-110 

acetoarsenite 109 

analysis  of,  notes 128 

arsonite 109 

compounds  with  cyanogen 107 

detection  of 109 

detection  of,  in  alloys 379 

detection  of  traces  of,  with  H2S .  .  .    108 

detection  of,  with  HBr 108 

electrical  conductors 104 

estimation  of 109 

ferrocyanide,  formation  of 107 

group,  metals  of 56,  100 

hydroxide  of 104 

occurrence  of 104 

oxides  of 104 

precipitation  of,  by  iron  wire 110 

preparation  of 104 

properties  of 104 

pyrites 204,  313 

reactions  of,   comparison  with  Bi 

andCd 112 

reduction  by  ignition 112 

reduction  of,  by  KCNS 107 

salts,     detection     by     potassium 

xanthate 107 

salts,     reaction     with     zinc-plati- 
num couple 110 

salts,  reduction  of,  with  H3PO2 ....   107 
salts,  separation  of,  from  Cd  by 

Na4P2O7 107 

salts,  solubilities  of 105 

separation    of,    from    Bi   by   gly- 

cerol 101 

separation   of,    from   Cd   by  gly- 

cerol 105 

separation  of,  from  Cd  by  Na2S2O3 

andNa^COs 112 

separation  of,  from  Cd  by  H2S  in 

presence  of  KCN 107 

separation    from    Cd   by   nitroso- 

/3-naphthol 107 

separation    from    Cd    by    ammo- 
nium benzoate 107 

separation  from  Pd 106 

traces,  loss  of 118 

traces  of,  with  K4Fe(CN)6 107 

Corundum 144 

Cream  of  tartar,  formation  of 260 


FAGE 

Crocoite 29,  148 

Crookesite 204 

Cryolite 144,  297 

Crythrite 57 

Cuprammonium  salts 106 

Cupric  hydroxide  in  NH4OH 105 

hydroxide,  effect  of  boiling 106 

hydroxide,  formation  of 106 

hydroxide,  with  glucose 106 

hydroxide,  with  tartrates 105,  106 

salts,  reaction  with  glucose 105 

salts,  reaction  with  iodides 108 

salts,  reaction  with  Na^Oa 108 

salts,  reduced  by  SO2 109 

sulphide,  colloidal 108 

sulphide,  formation  of 108 

sulphide,   separation  from  Cd  by 

H2SO4 108 

sulphide,  solubility  in  (NH4)2SX ....  108 

sulphide,  solubility  in  KCN 108 

sulphide,  with  K2S 118 

sulphide,  with  (NH4)2Sz 115 

Cuprite 104 

Cuprous  iodide 108 

oxide,  formation  of,  by  glucose. .  .  .  105 

salts,  oxidation  of,  by  As2O3 110 

salts,  separation,  from  Cd  by  S. .  .  107 

salts,  with  metallic  sulphides 108 

sulphide,  formation  by  NaoS2O3. .  .  108 

thiocyanate,  formation  of 107 

Cyanates,  detection  of,  in  presence 

of  cyanides 279 

Cyanic  acid 279 

Cyanide  of  silver,  distinction  from 

chloride 273 

Cyanides,  detection  as  thiocyanate. .  275 

double,  dissociated  by  acids 272 

double,  not  dissociated  by  acids. .  .  272 

estimation  of 275 

guaiacum  test 275 

ignition  of 274 

preparation  of 272 

reactions  with 272 

simple,  with  mineral  acids 273 

solubility  of 272 

transposition  by  acids 275 

Cyanogen     properties     and     reac- 
tions                                      .  271 


Danger  and  Flandin,   detection    of 

arsenic 69 

Daubreelite 148 

Decomposition    of    organic    mate- 
rial  374-375 


426 


INDEX. 


Dialysis,    separation    from    organic 

material  by 375 

Diamond 254 

Diaspore 144 

Didymium 199 

Didymium  Earths 202 

Dimethylaniline,    test    for     nitric 

acid 290 

Dimethylglyoxime,  test  for  nickel  176 
Diphenylamine,  test  for  nitric  acid  290 

Dissociation,  electrolytic 20 

Dithionic      acid,      formation      and 

properties 324 

Dolomite 220,  267 

Dragendorff's  reagent 102 

Electrolytic  dissociation 20 

Enargite 57 

Epsom  salts 220,  313,  331 

Equations      illustrating      oxidation 

and  reduction 409 

rule  for  balancing 246 

Erbium 200 

Ethyl  acetate,  odor  of 257 

Euxenite 137,  198,  202,  207 

Everett's  salt 157 

fatty  material,  removal  of 375 

Feldspar 144 

Ferric  acetate,  formation  of 257 

acetate,  separation  of  from  chro- 
mium   157 

basic     nitrate,     separation     from 

aluminum 161 

and   ferrous    compounds,    distinc- 
tion    165 

hydroxide,  antidote  for  arsenic ....  62 

phosphate,  formation  of 159 

salts,  detection  of  traces 158 

salts,  with  acetates 157 

salts,  with  BaCOa 156 

salts,  with  HI  and  iodides 161 

salts,  with  H3PO2 159 

salts,  with  H2S 150 

salts,  with  KCNS 158 

salts,  with  K3Fe(CN)6 158 

salts,  with  K4Fe(CN)6 157-158 

salts,  with  stannous  chloride 89 

salts,  separation  from  ferrous  sul- 
phate   156 

Ferric     thiocyanate,      distinction 

from  ferric  acetate 157 

hindrance  to  reactions  of 158 

Ferricyanides,     in    distinction    be- 
tween Co  and  Ni. .  .  .170 


Ferricyanides,  reactions  of 278 

Ferrite 154 

Ferrocyanides,  detection  of 277 

detection  and  estimation 279 

reactions  of 276-277 

Ferrotellurite 138 

Ferrous  iron,  detection  of,  in  ferric 

salts 158 

in  the  third  group 164 

in    the    third    group    with    phos- 
phates     194 

salts,  traces  in  ferric  salts 158 

salts,  with  chromates 161 

salts,  with  HNO3 159 

salts,  with  KCN 157 

Ferrous  salts,  with  K3Fe(CN)6 158 

salts,  with  K4Fe(CN)6 157 

sulphate,  with  gold  salts 93 

First  group  metals,  table  of 52 

Fixed  alkalies 227 

alkali  hydroxides  on  stibine 79 

alkalis  with  salts  of  tin 84 

Flame,  blowpipe,  production  of 376 

or  color  tests 385 

oxidizing  and  reducing 375-376 

reactions  with  copper  salts .  .   109 

Flint 299 

Fluor-spar 216,  297 

Fluorides,  solubilities  of 298 

Fluorine 297-298 

Fluosilicates,  formation  of 298-299 

Fluosilicic  acid 298-299 

in  detection  of  potassium 231 

in  separation  of  Ba,  Sr  and  Ca .  .  .  213 
Formates,     formation    from    cyan- 
ides     274 

Fourth  group,  directions  for  anal- 
ysis     189 

reagents 142 

sulphides  colloidal 189 

table  of 183 

Fresenius   and   Babo,   detection  of 

arsenic 68 

Froehde's  reagent 99 

Fulminating  gold 92 

Gadolinite 202,  207 

Galena 29,  45,  313 

Gallium  (eka-aluminium) 200 

Gamma  iron 154 

Garnierite 173 

Gas-laws 21 

Gases,  absorption  of  by  palladium ...  133 
Germanium,    properties   and  reac- 
tions. .  .    137 


INDEX. 


427 


Germanium  sulphide 118 

Glass,  etching  by  hydrofluoric  acid.   313 

Glauber's  salts 313,  331 

Glucinum  (Beryllium) 200-201 

distinction  from  yttrium 207 

separation  from  aluminum 201 

separation  from  cerium 198 

Glucose,    in    formation   of   cuprous 

oxide 105 

Gold 91-93 

detection  in  alloys 379,  380 

detection  of 93 

distinction  from  Pd 133 

estimation  of 93 

fulminating 92 

notes  on  analysis 125 

occurrence,  properties,  etc 91 

reduction  by  ferrous  sulphate 93 

reduction  with  oxalic  acid 92 

salts  with  alkalis 92 

salts  with  stannous  chloride 89 

separation  from  Ir 135 

Graphite 254 

Greenockite 101,  110 

Gypsum 216,  219,  313,  331 

Haematite 155 

Halogens 9 

as  oxidizers 340 

compounds,  comparative  table  of  373 

hydracids  as  reducers 340 

separation     of      by     persulphate 

method 347,  403 

Hausmannite 177 

Heat,     upon    substances    in    closed 

tubes 376,382 

upon   substances   in   open   tubes. 

376,  383 

Heavy  spar 313,  331 

Hydriodic  acid 365-368 

action  on  antimonic  salts 78 

action  on  arsenic  salts 61 

action  on  ferric  salts 161 

as  a  reducer 366,  367 

formation  of 365 

Hydrobromic  acid 356-360 

detection  of  Cu  with 108 

formation  of 357 

occurrence  of 357 

preparation  of 357 

properties  of 357 

•     reactions  of 357 

Hydrochloric  acid 341-348 

action  on  Sb-jSn 77 

action  on  bismuth  nitrate 102 


Hydrochloric  arid,  effect  of  excess 

in  second  group 114 

formation  of 341 

formation  from  MgCl2 222 

gas  on  arsenic  sulphide 67 

occurrence  of 341 

preparation  of 342 

properties  of 341 

reactions  with 343 

solubilities  of 342 

Hydrocyanic  acid 271-275 

formation  of 272 

occurrence  of 272 

on  PbO2 274 

preparation  of 272 

properties  of 271 

solubilities  of 272 

Hydrof erricyanic  acid 277-279 

Hydroferrocyanic  acid 275-277 

separation  from  hydroferri- 

cyanic  acid 277 

Hydrofluoric  acid 298 

Hydrofluosilicic  acid  (fluosilicic 

acid) 298 

Hydrogen 250,  251 

absorption  by  Pd  sponge 13& 

detection  of 251 

estimation  of 251 

formation  of 250 

nascent 251 

occluded 251 

occurrence  of 250 

preparation  of 250 

properties  of 250 

reactions  with 250 

reducing  action  of,  with  ignition.  . .  251 

solubilities  of 250 

peroxide,  detection  of 296 

peroxide,  estimation  of 297 

peroxide,  estimation  of  bismuth 

with 104 

peroxide,  formation  of 295 

peroxide,  occurrence  of 295 

peroxide,  on  sulphides  of  arsenic 

and  antimony 121 

peroxide,  preparation  of 295 

peroxide,  properties  of 294 

peroxide,  reactions  with 295 

peroxide,  reagent  to  separate  Co 

from  Ni 190 

peroxide,  separation  from  ozone. .  .  243 
peroxide,  separation  of  Al,  Fe  and 

Cr  with 152 

peroxide,  solubilities  of 295 

peroxide,  with  arsenic 71 


428 


INDEX. 


Hydrosulphuric  acid 315-320 

action  on  copper  salts 108 

action  on  ferric  salts 160 

dissociation  of 114 

formation  of 316 

gas  as  a  reagent 113 

gas  on  antimony 67 

gas  on  arsenic 67 

occurrence  of 316 

on  aluminum  salts 146 

on  stannic  salts 86 

on  stannous  salts 85 

on  third  and  fourth  group  salts .  142,  164 

preparation  of 316 

properties  of 315 

uses  as  a  reagent 317 

with  arsenic  acid 114 

with  oxidizing  agents 114 

Hydrosulphurous  acid 323,  324 

Hydroxylamine,      formation      and 

properties 286 

Hydrozoic  acid 8,  282 

Hypobromous  acid,  formation  and 

properties 360 

Hypochlorites,  detection  of 404 

formation  of 348 

formation  from  chlorine 344 

on  arsenic 66 

Hypochlorous  acid 348,  349 

Hypoiodous  acid,  existence  of 363 

Hyposulphites,  detection  of 305 

ignition  of 305 

Hypophosphites    in    formation    of 

PHS 305 

Hypophosphoric  acid 307 

Hypophosphorous  acid 304-306 

estimation  of 306 

formation  of 304 

preparation  of 305 

properties  of 304 

reactions  of 305 

solubilities  of 305 

with  bismuth  salts 102 

Hyposulphurous  acid 323-324 


Imperial  green 109 

Indigo  test  for  nitric  acid 289 

Indium 201 

Ink,  common 157 

sympathetic 168 

lodates   detection  of 371 

estimation  of 371 

formation  of 369 

ignition  of 371 


lodates,  reactions  of 370 

lodicacid 369-371 

formation  of 369 

preparation  of 369 

properties  of 369 

reactions  of 370 

Iodide  of  nitrogen 363 

Iodides,  decomposition  by  HNO3 . .  .   290 

detection  as  PdI2 133 

detection  of 368 

estimation  of 369 

formation  of 365 

ignition  of 368 

occurrence  of 365 

reactions  of 366 

separation  of,  from  bromides  and 

chlorides  by  KMnO4 181 

solubilities  of 365 

Iodine 362-364 

detection  of 364 

estimation  of 364 

formation  of 363 

liberation  by  copper  salts 108 

occurrence  of 363 

on  antimonous  salts 78 

o"n  antimony 66 

on  arsenic 66 

preparation  of 364 

properties  of 362 

reactions  of 363 

separation  from  Br  by  Pd 134 

solubilities  of 363 

Ions 21 

lonization  and  solution 20-24 

Iridium 134,  135 

Iron 157-162 

alpha 154 

and  zinc  groups 141 

beta 154 

detection  of 165,  166 

detection  of  traces  in  copper 157 

detection  of  traces 157,  158 

estimation  of 162 

gamma 154 

group 144 

group,    separation   from    Co,    Ni, 

and  Mn  by  ZnO 161 

hydroxides 155 

in  relation  to  metals 6 

occurrence  of 154 

oxidation  of 162 

oxides, 154,  155 

preparation  of 154 

properties  of 153 

pyrites 313 


INDEX. 


429 


Iron,  reduction 162 

salts,  ignition  of 161 

salts,  solubilities  of 156 

salts,  with  alkalis 156 

salts,  with  nitroso-/3-naphthol 157 

salts,  separation  from  Al  as  basic 

nitrate 161 

separation  from  Al  and  Cr  by 

nitroso-/3-naphthol 157 

separation  from  Cr  and  Al 157 

separation  from  Ni  by  xanthate ...  175 

solubilities  of 155 


Kainite 228 

Kieserite..  .  313 


Lanthanum 202 

Lead 29-36 

acetate,  properties  of 32 

chloride 34 

chloride,  precipitation  of 53 

chromate,  formation  of 35 

compounds,  ignition  of 35 

detection  in  alloys 379 

detection  of 36 

estimation  of 36 

in  the  test  f or  Al 167 

iodide,     formation     and     proper- 
ties   35 

notes  on  analysis  of 129 

occurrence  of 29 

oxidation  and  reduction 36 

oxides  of 29 

oxides,  solubilities  of 30 

preparation  of 29 

properties  of 29 

red 29 

relation  to  nitrogen  family 7 

salts,  reactions 32-35 

salts,  solubilities  of 31 

solubilities  of  metallic 30 

sulphate,    formation    and   proper- 
ties of 34 

sulphide,    formation    and    proper- 
ties for 33 

tests  for 54 

Leblanc-soda  process 259 

Lepidolite 241 

Light,  action  on  silver  salts 50 

Lime,  slacked 217 

Limestone  (CaCO3) 216,  219,  267 

Lithium 234-236 

Limonite 155 


Linnaeite 167 

Lollingite 57 

Magnesia  mixture. . . 146 

Magnesite 220,  267 

Magnesium 220-222 

as  a  reducing  agent 222 

detection  of 222 

estimation  of 222 

hydroxide,  formation 220 

occurrence  of 220 

oxalata,    separation    of,    from    K 

andNa 221 

oxide,  formation  of 220 

preparation  of 220 

properties  of 220 

removal  for  detection  of  sodium. .  .  242 

salts,  with  ammonium  salts 221 

salts,  with  arsenic  acid 61 

salts,  with  Na2S 221 

salts,  solubilities  of 220 

Magnetite 155 

Malachite 104 

Manganates,  identification 405 

Manganese 177-182 

detection  of,. . . . '. 182,  191 

estimation  of 182 

hydroxides  of 177 

hydroxides,  solubilities  of 178 

ignition  of 182 

in  third  group 164,  166-167,  189 

occurrence  of 177 

oxidation  of 182 

oxidation  to  permanganic  acid ....   180 

oxides 177 

oxides,  solubilities  of 178 

preparation  and  properties 177 

reduction  of 182,  183 

reduction  by  sulphites 181 

salts,  reactions  with  oxalic  acid .  .  .    180 

salts,  solubilities  of 178 

salts,  with  alkalis 179 

salts,  with  sulphides 181 

separation  from  zinc  with  acetic 

acid 189 

solubilities  of 178 

with  KI 181 

Manganic  acid 177 

Magnesite 177 

Marble 216,267 

Marcosite 155 

Marsh's  test 62 

Mass  act  ion,  law  of 23,  38 

Mayer's  reagent 43,  238 

Melanconite .104 


430 


INDEX. 


Mercurammonium  compounds ...     39 
Mercuric    chloride    with    stannous 

chloride 88 

sulphide,    formation    and    proper- 
ties      41 

sulphide,  with  K2S 118 

Mercury 37-45 

chlorides 42 

compounds,  ignition  of 43 

detection  and  estimation  of 44 

iodides 42 

metallic,  analysis  of 379 

occurrence  of 37 

oxidation  of 45 

oxides 37 

preparation  and  properties  of 37 

salts,  reactions 39,  43 

salts,  solubilities  of 38 

solubilities  of 37 

sulphide,  analysis  of 128 

Metals,  classification 10 

grouping 387 

table  of  separation 388 

Metaphosphoric  acid 308 

Metastannic  acid 83 

Mica 299 

Microcosmic  salt 236 

use  in  ignition 377 

Milk  of  lime 217 

Millerite 173 

Molybdates  in  analysis 54 

with  phosphates 98 

Molybdenite 97 

Molybdenum 97-99 

deportment  in  second  group 99 

detection  of 99,  124 

estimation  of 99 

ignition  tests 99 

notes  on  analysis  of 124 

occurrence  of 97 

oxides  and  hydroxides 97 

preparation  and  properties 97 

reduction  tests 99 

solubilities  of 97 

Molybdic  acid 97 

Molybdite 97 

Mottramite 206 

Monazite 204 

Monazite  sand .  .  .199 


Nascent  hydrogen  on  nitric  acid. .  .  .   286 

Neodymium 199,  203 

Nessler's  reagent 43,  237 

Niccolite 57,  173 


Nickel 173-176 

detection  of 176 

detection    of,    in   presence   of    Co 

by  KI 190 

distinction  from  cobalt 175 

estimation  of 176 

hydroxides 173 

hydroxides  with  KI 176 

ignition  of 175 

occurrence  of 173 

oxidation  of 177 

oxides 173 

properties  and  preparation 173 

reduction 177 

salts  with  alkalis 174 

separation      from      Co,      cyanide 

method 169-170 

separation   from   Co,    by  nitroso- 

j8-naphthol 173 

separation    from    Co,    by    KNO2 

170-171 

separation  from  Co,  by  sulphide  175 
separation     from     Co,     by     xan- 

thate 175 

separation  from  cobalt  by  NH4CNS  172 

solubilities  of 173 

solubility    of    NiS    in   ammonium 

sulphide 175 

test  for  by  means  of  dimethylglyox- 

ime 176 

xanthate,  separation  from  Fe 175 

Niobium  (Columbium) 198-199 

Nitrates,     decomposition    by    igni- 
tion    288 

distinction  from  chlorates 404 

occurrence  of 285 

preparation  of 285 

proof  of  absence 402 

solubilities  of 286 

Nitric  acid 285-291 

as  an  oxidizer 286 

Brown  ring  test 288 

decomposition  of,  by  HC1 287 

detection  of 288 

detection  by  diphenylamine 289 

detection  by  reduction  to  NH3..286,  289 
detection  by  reduction  to  nitrite .  .   289 

dissociation  by  heat 288 

estimation  of 291 

formation  of 285 

indigo  test 289 

in  separation  of  Sn,  Sb  and  As .  . .   121 

sodium  salicylate  test 289 

test   for   by    dimethylaniline    and 

diphenylamine 290 


INDEX. 


431 


Nitric  acid,  with  phenol 290 

with  pyrogallol 291 

with  brucine 290 

occurrence  of 285 

on  antimony 66 

on  arsenic 66 

preparation  of 285 

products  of  reduction 286 

properties  of 285 

Nitric  anhydride,  formation  of 286 

oxide 104,  221,  283 

Nitrites,  decomposition  by  ignition.   284 

detection  of 284 

test  for  nitric  acid 288-290 

Nitrof erricyanides 278 

Nitrogen 281,  282 

chloride 62,  120,337 

combination  with  elements 282 

detection  and  estimation 282 

family 7 

formation,  occurrence 282 

peroxide 285 

properties 281 

Nitroso-/3-naphthol,   separation  of 

CoandNi 170,  190 

separation  of  Cu  from  Cd 107 

with  iron  salts 157 

Nitroprussides 278 

Nitrous  acid 284-285 

as  an  oxidizer. .  •. 284 

as  a  reducer 284 

formation  of 284 

occurrence  of 284 

properties  of 284 

reactions  with 284 

solubilities  of 284 

Noble  metals,  enumeration 7 

Nordhausen  sulphuric  acid 332 

Notes  on  detection  of  acids 401 

on  analysis  of  calcium  group. .  .224—226 
on  analysis  of  third  group 164 

Opal 299 

Orangite 204 

Order  of  laboratory  study 25 

Organic  substances,  removal  of .  374,  375 

Orpiment 57,  313 

Orthoclase 144 

Osmium 133 

Osmotic  pressure 20,  21 

Oxalates,  decomposition  by  ignition 

of 402 

decomposition  by  oxidation 402 

detection  of 266 

distinction  from  tartrates 260,  389 


Oxalates,  estimation  of 266 

ignition  of 266 

in  3d,  4th  and  5th  groups 194 

reactions  of 264 

solubilities  of 264 

Oxalic  acid 263-266 

as  a  reducer 264 

decomposition  of  by  H.SO4 265 

formation  of 263 

in  separation  of  gold 92 

occurrence  of 263 

preparation  and  properties  of ^263 

solubility  of 264 

Oxidation,    balancing   equations   in 

244,  245 

Oxidizing  flame 375 

Oxygen 291-293 

as  a  poison 292,  293 

combinations  with  ignition 293 

detection  of 293 

estimation  of 293 

foriuLiuan  of 291 

occurrence  of 291 

preparation  of 292 

reactions  with 292 

Ozone 293,294 

separation  from  H2Q2 295 


Palladium 133, 134 

distinction   from   gold   and   plati- 
num    134 

separation  from  copper 106 

sponge 133 

Palladous  iodide  in  analysis 133 

Paris  green 62,  109 

Pearlite 154 

Pentathionic   acid,   formation  and 

properties ' 325 

Pentlandite 173 

Perchlorates,       preparation       and 

properties 353,  354 

Perchromic  acid 153 

Periodic  acid 372 

system,  table  of 2 

Permanganates  identification 405 

action  on  antimonous  salts 78 

Permanganic  acid 178 

Persulphate    method  of  separating 

the  halogens 347,  403 

Persulphuric  acid 336 

Petalite 241 

Phenol  reaction  for  nitric  acid 290 

Phenylhydrazine,      on     aluminum 

salts .145 


432 


INDEX. 


Phosgene,  formation 262 

Phosphates,  changes  by  ignition. . .  312 

detection 165,  312,  402 

distinction       between       primary, 

secondary  and  tertiary 310 

estimation  of 313 

in  presence   of  third  and  fourth 
group  metals.. 143,  193,  194,  196,  197 

occurrence  of 308 

reaction  with  ammonium  molyb- 

date 193,311 

separation  as  ferric  phosphate 193 

solubilities  of 309 

Phosphides,  formation  of 312 

Phosphine 304 

Phosphoric  acid 307-313 

preparation  of 309 

properties  of 307 

Phosphoric    anhydride,    formation 

of 308 

Phosphorite 216 

Phosphorous  acid 306,  307 

detection  of 307 

preparation  and  properties  of 306 

Phosphorus 301-304 

detection  and  estimation  of 304 

in  combination  with  the  halogens  304 
occurrence  and  preparation  of ....   303 

properties  of 301,  303 

use  in  match-making 302 

Phosphotungstates 137 

Picric   acid,    in  detection  of  potas- 
sium   230 

Pitch-blende 206 

Plaster  of  Paris  (calcium  sulphate). .  219 

Platinized  asbestos 93 

Platinum 93-97 

apparatus,  care  of 95 

black 93 

chloride,  as  a  reagent 94 

detection  of 96,  124,  379-380 

distinction  from  palladium 134 

estimation  of 96 

iridium  alloys,  properties 134 

notes  on  the  analysis  of 125 

occurrence  of 94 

preparation  and  properties 93,  94 

reduction  of 95,  96 

sponge 93 

Polarity 3 

Potassium 228-232 

as  a  reducing  agent 232 

bichromate,     in    test    for    stron- 
tium and  calcium 225 

carbonate,  as  a  reagent 229 


PAGE 

Potassium  chlorate,  in  preparation 

of  oxygen 292 

chloride  with  platinum  chloride ...     95 

cyanide  with  copper  salts 107 

cyanide  with  ferrous  salts 157 

detection  of 229,  232 

estimation  of 232 

ferricyanide,  formation  of 277 

ferrocyanide,  formation  of 273,  276 

hydroxide,  as  a  reagent 229 

iodate,   in   separation   of  alkaline 

earths 213 

iodide,  as  a  reagent 240 

iodide,    in    separation    of     AgCl 

from  SbCl3 122 

iodide,  in  the  test  for  nickel 190 

iodide,  on  nickelic  hydroxide 175 

iodide,  on  permanganates 181 

nitrite    in    separation    of    cobalt 

from  nickel 170, 171 

occurrence,        preparation        and 

properties  of 228 

picrate 230 

pyroantimonate 73,  224 

salts,  flame  test 231 

thiocyanate  with  copper  salts 107 

thiocyanate  with  iron  salts 158 

xanthate,  for  detection  of  copper  107 

Powder  of  algaroth 75 

Praseodymium 199,  202 

Precipitates,     formation     and     re- 
moval of 17,  18 

Principles 405 

Problems     in     molecular     propor- 
tions       19 

in  synthesis 410 

Proustite 46,  57 

Prussian-blue,  formation  of .  . .  158,  274 

Purple  of  Cassius 89,  93 

Pyrargyrite 46 

Pyrite 45,  155 

Pyroantimonic  acid 73 

Pyrochlor 198 

Pyrogallol,      as    a    test    for    nitric 

acid 291 

Pyrolusite 177 

Pyromorphite 29 

Pyrophosphoric  acid,  formation. . . .  308 

Pyrosulphuric  acid,  formation 332 

Pyrrhotite 155,  173 

Quartz 299 

Reagents,  care  in  the  addition  of . . .     17 
list  of...  .  415 


INDEX. 


433 


Realgar 57,  313 

Reducing  flame,  description  of 375 

Reduction,   balancing  equation"  in 

244,  245 

with  charcoal 376,  377,  381 

Reinsch's  test  for  arsenic 67 

Rhodium,  distinction  from  iridium  135 

properties  and  reactions 132 

Rhodocrosite 177 

Rochelle  salts,  composition  of 260 

Rosolic  acid  as    a  test  for  carbon 

dioxide 270 

Rubidium,     properties    and    reac- 
tions    240 

Ruby 144 

Rule  for  balancing  equations 246 

Ruthenium,    properties   and   reac- 
tions     129 

Rutile 205 

Salt 233 

Saltpeter,  occurrence 285 

Samarium,  properties  and  reactions  202 

Sand 299 

Sapphire 144 

Scandium,    properties     and     reac- 
tions  202-203 

Scheele's   green   and  Schweinfurt's 

green 109 

Scheelite 136 

Selenic   acid,   separation  from  sul- 
phuric acid 140 

Selenite 216 

Selenium,     properties     and     reac- 
tions     139,  140 

Siderite 155 

Silica  (silicon  dioxide) 300 

detection  and  estimation  of 301 

in  the  microcosmic  bead 301 

in  the  third  group 167 

removal  of 402 

solubilities  of 300 

Silicates,    decomposition    by    igni- 
tion   300 

in  analysis 54 

Silicic  acid 299-301 

Silicon 299 

distinction  from  tantalum 203 

Silico-fluoride  (fluosilicate) 298 

Silicon  fluoride,  formation 297,  298 

preparation  and  properties 298 

separation  from  thorium 205 

Silver 45-50 

arsenate  and  arsenite,  formation . .     62 
bromate,  properties  of 361 


Silver  chloride,  formation  and  prop- 
erties      48 

cyanate      in      distinction       from 

cyanides 279 

detection  of 50,  380 

estimation  of 50 

in  presence  of  mercury  salts 55 

iodate,  properties  of 370 

mirror,  formation  by  tartrates ....  261 

nitrate,  action  on  stibine 79 

nitrate   with   stannous   and   anti- 

monous  salts 78,  79,  88 

occurrence  and  properties  of 45 

salts,  action  of  light  upon 50 

solubilities  of 46 

thiocyanate,  separation  from  silver 

chloride 281 

Sipylite 207 

Skutterudite 167 

Smaltite 57,  167 

Soapstone 299 

Soda  lime  on  stibine 79 

process,  Le  Blanc's 267 

•     process,  Solvay's 268 

Sodium 232^235 

amalgam,  action  with  arsenic 64 

as  a  reducing  agent 235 

detection  of 73,  234 

estimation  of 235 

flame  test 234 

hydroxide,  formation  of 233 

nitroferricyanide  as  reagent. .  .   236,  320 

occurrence  of 233 

phosphate  as  reagent 233 

Sodium   phosphomolybdate    as  re- 
agent..  98,  238 

preparation  and  properties  of.  .232,  233 

pyroantimonate 73,  80 

'  pyrophosphate   with    copper   and 

cadmium 107 

salicylate  test  for  nitric  acid 289 

sulphide,  preparation  of 317 

thiosulphate  on  cupric  salts 108 

thiosulphate  with  antimony  salts. .     78 
thiosulphate     with     third     group 

metals 146,  147 

Solids,  conversion  into  liquids 328 

decomposition  upon  ignition. .  .  382,  383 
effect    on    ignition    with     cobalt 

nitrate 384 

preliminary  examination  of 375 

separation  of 17 

table    for    preliminary    examina- 
tion    382 

Solubility,  degrees  of 15,  16 


434 


INDEX. 


Solubility-product 24 

Solution  and  ionization 20-24 

Solvay  soda  process 268 

Sonnenschein's  reagent 98 

Sperrylite 94 

Stannic  salts,  solubilities 84 

sulphide,    formation    and    proper- 
ties of 86 

Stannite 82 

alkali,  as  a  test  for  bismuth 103 

Stannous  chloride  on  mercury  salts  43 

chloride  as  a  reducing  agent 88 

chloride  with  gold  salts 93 

chloride  with  molybdic  acid 99 

salts,     distinction     from     stannic 

salts 125 

salts,  solubilities 84 

salts  with  silver  nitrate 87 

salts  with  sulphurous  acid 86 

sulphide,    formation    and    proper- 
ties   85 

Stephanite 46 

Stibine,     decomposition     by     soda 

lime 79 

formation  of 79 

reaction    with    fixed    alkali    hy- 
droxides    79 

reaction  with  silver  nitrate 79 

separation  from  arsine 65 

Stibnite 72 

Strontianite 214 

Strontium 214-216 

detection  of 216,  219 

estimation  of 216 

hydroxide,  formation 214 

occurrence  of 214 

preparation  and  properties  of 214 

sulphate,  distinction  from   CaSO4  215 

sulphate,  separation  from   BaSO4  215 

Sulphates,  detection  of 335 

estimation  of 336 

ignition  of 335 

preparation  of 332 

reduction    by   ignition  with    car- 
bon    256 

solubilities  of 333 

Sulphites,  detection  of 330 

distinction  from  sulphates ,  .  330 

estimation  of 330 

ignition  of 330 

interference  in  test  for  oxalates .  .  .  402 

preparation  of 327 

separation  from  sulphates  by  Ba 

salts 213 

solubilities  of .  328 


Sulphides,  detection 320 

estimation  of 321 

formation  of 316 

ignition  of 319 

reactions  of 318,  319 

solubilities  of 28,  317 

Sulphur 313,  315 

combinations  on  ignition  of 315 

detection  and  estimation  of 315 

formation  of 313 

in  the  tin  group 118 

occurrence  of 313 

oxidation  by  reagents 314,  315 

oxides 313 

precipitation  of 54,  114,  115 

preparation  and  properties  of  .  .313,  314 
reactions  in  forming  sulphides.  .  .  .  314 

relations  of 9 

separating  copper  from  cadmium  107 
solubilities  of 314 

Sulphuric  acid 331-336 

detection  in  presence  of  sulphates  335 

formation  and  occurrence  of 331 

manufacture  of 331 

properties  of 331 

reactions  with 333-335 

separation  from  Se 140 

separation  from  Te 139 

anhydride,  preparation  of 331 

Sulphurous  acid 327-330 

on  arsenic  acid 60 

and  sulphites  as  reducers 329 

occurrence  of 327 

preparation  and  properties  of 327 

formation  of 327 

reduction  of  cupric  salts 108 

solubilities  of 328 

on  stannous  salts 86 

Sylvanite 138 

Synthesis,  problems  in 410 


Table  for  acids  as  precipitated  by 

barium  and  calcium  chlorides . . .  398 

for  acids  precipitated  by  silver 
nitrate 399 

for  acids,  preliminary 390 

for  analysis  in  presence  of  phos- 
phates by  the  use  of  alkali  ace- 
tates and  ferric  chloride 196 

for  analysis  in  presence  of  phos- 
phates by  use  of  ferric  chloride 
and  barium  carbonate 107 

for  analysis  of  the  Silver  Group 
(first) 52 


INDEX. 


435 


Table    for  analysis  of  the   Copper 

Group  (second) 126 

for    analysis    of    the    Tin    Group 

(second) 116 

for    analysis   of   the    Iron    Group 

(third) 163 

for   analysis   of   the   Zinc    Group 

(fourth) 188 

for  analysis  of  the  Calcium  Group 

(fifth) 223 

of  grouping  of  the  metals 387 

of  separations  of  the  metals 388 

of    separation   of   the   ammonium 
sulphide     precipitates     of     the 

Iron  and  Zinc  Groups 192 

of  solubilities 411 

Tannic  acid  with  iron  salts 157 

Tantalite 198,  203 

Tantalum,  distinction  from  silica  .  .   203 

distinction  from  titanium 203 

properties  and  reactions  of 203 

separation  from  columbium 203 

Tartar  emetic,  composition  of 260 

Tartaric  acid 259-261 

in  detection  of  potassium 229 

distinction  from  citric  acid 259 

formation  and  properties 259-260 

Tartrate  calcium,  deportment  with 

water 260 

detection  of 260 

distinction  from  citrates 260 

distinction  from  oxalates 260,  401 

estimation  of 261 

Tartrates,  ignition 261 

reactions 260 

solubilities 260 

Tellurite 138 

Tellurium 138-139 

distinction  from  selenium ....  139,  140 
properties  and  reactions  of ....  138,  139 
separation  from  sulphuric  acid.  . .  .  139 

Tenorite 104 

Terbia 203 

Terbium 203-204 

Tetradynite 138 

Tetrathionic   acid,   formation    and 

properties 325 

Thallious  iodide 204 

Thallium,  properties  and  reactions .  204 
Thioacetate   in   formation   of    sul- 
phides   316 

Thiocyanates,  reactions  with 280 

test  for  cobalt  by  means  of 170 

Thiocyanic  acid  as  a  reducer 281 

properties  of 280 


Thionic    acids,    table    of    compari- 
sons     326 

Thiosulphates,  detection  of 323 

distinction    from    sulphates     and 

sulphites 323 

estimation  of 323 

ignition  of 323 

formation  and  properties  of 321 

Thiosulphuric  acid 321-323 

Third  group  reagents 142 

Thorite 204 

Thorium 204-205 

Tin 82-89 

creaking  of 82 

detection  of 88,  122,  379 

estimation  of 88 

Group,  metals  of 56 

Group,    separation    from  Copper 

Group 115 

Group,  sulphides  with  (NH4)2S.r. .  .   115 

occurrence  of 82 

oxidation  of 88 

oxides  and  hydroxides 82 

preparation  and  properties  of 82 

notes  on  the  analysis  of 124,  125 

relation  to  Nitrogen  Family 7 

reduction  by  ignition 87 

salts  with  the  alkalis 84 

salts  with  hydrosulphuric  acid ....     85 

separation  from  antimony 81 

separation    from     antimony    sul- 
phides      123 

.   separation  from  arsenic 118 

solubilities  of 83 

sulphides,  colloidal 115 

with  antimony  and  with  arsenic  .  .     87 

Tinstone 82 

Titanif erous  iron 205 

Titanite 205 

Titanium 205-206 

distinction  from  columbium 206 

distinction  from  tantalum 203 

properties  and  reactions  of 205 

separation  from  thorium 205 

Triphylite 241 

Trithionic     acid,     formation     and 

properties 324 

Tungsten,  properties  and  reactions 

136, 137 
Turnbuii's  blue. .  .   155 


Unit  of  quantity 23 

Uranium,  properties  and  reactions.   206 
Urea,  from  ammonium  cyanate 279 


436 


INDEX. 


Valence,  negative 3 

Valentinite 72 

Vanadinite 206 

Vanadium 206, 207 

Volatile      alkali      (ammonium     hy- 
droxide)    221 

Volborthite.  .  ....  206 


Wad 167 

Water,  action  on  bismuth  salts 101 

action  on  antimonous  salts 75 

Welsbach  burners 209 

Witherite 211 

Wolframium  (tungsten) 136 

Wulfenite .  .  97 


Ytterbium    properties    and    reac- 
tions   207 

Yttrium..  .  207 


PAGE 

Zincates,  formation  of 184 

Zinc 183-186 

blende: 183,  313 

blende  (Freiburg) 201 

detection  and  estimation  of 186 

Family 5 

granulated 63,  183 

Group,  table  for  analysis 188 

Group,  comparative  reactions 187 

hydroxide  and  oxide 184 

ignition  of 186 

occurrence  of .  . 183 

oxidation  of 186 

platinized 183,  250 

preparation  and  properties 183 

reduction  of 186 

salts,  solubilities  and  reactions  of  184 
sulphide,    formation    in    presence 

of  acetic  acid 184 

Zircon 209 

Zirconium  .  ,  .209 


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